PRACTICAL  AND  ANALYTICAL 


CHEMISTRY. 


BEING 


A  COMPLETE  COURSE  IN  CHEMICAL 
ANALYSIS. 


BY 


HENRY   TRIMBLE,   Pn.G., 

PROFESSOR    OF   ANALYTICAL   CHEMISTRY    IN   THE    PHILADELPHIA   COLLEGE 
OF   PHARMACY. 


OF  TOT! 

UNIVERSITY 


SECOND  EDITION. 

REVISED,  ENLARGED  AND  ILLUSTRATED. 


PHILADELPHIA: 

P.    BLAKISTON,    SON   &   CO., 

No.  1012  WALNUT  STREET. 
1886. 


COPYRIGHT,  1886,  by  P.  BLAKISTON,  SON  &  CO. 


PREFACE  TO  FIRST  EDITION. 


The  increased  amount  of  time  devoted  by  students  of  phar- 
macy and  medicine  to  analytical  chemistry  has  directed  more 
attention  to  the  subject  of  imparting  instruction  in  this  science. 

The  object  of  the  present  volume  is  to  place  before  the 
student  as  compact  a  course  as  possible,  in  order  to  enable 
him  to  become  familiar  with  the  subject  in  the  necessarily 
limited  time  at  his  disposal. 

The  author's  experience  has  led  him  to  believe  that  a  study 
of  Qualitative  Analysis  should  be  preceded  by  some  experi- 
ence in  the  preparation  of  the  more  important  gases  and  a  few 
of  the  salts.  Such  practice  requires  the  student  to  familiarize 
himself  with  the  construction  of  apparatus  as  well  as  with  the 
processes  of  filtration,  evaporation,  crystallization,  ignition,  etc. 

The  examples  for  preparation  may  be  increased  at  the 
option  of  the  instructor. 

In  Part  II  the  student  should  perform  the  reactions  of  each 
group,  and  then  be  furnished  with  a  solution  containing  some, 
or  all,  of  the  bases  of  the  group.  This  should  be  followed  by 
a  solution  in  which  he  should  search  for  all  the  elements 
previously  studied ;  such  practice  being  repeated  until  he  can 
correctly  determine  all  the  bases  present. 

In  order  to  enable  the  student  to  see  the  comparative  effect 
of  the  group  reagents  readily,  a  summary  has  been  introduced 
at  the  end  of  each  group.  This  he  should  be  able  to  write 


VII 


viii  PREFACE   TO    THE   FIRST    EDITION. 

out,  without  the  use  of  the  book,  before  attempting  to  analyze 
a  group  solution.  By  progressive  steps  he  is  thus  led  rapidly 
on  to  successfully  examine  the  more  complex  solutions  for 
both  bases  and  acids. 

The  grouping  of  the  bases  is,  to  a  certain  extent,  new,  but  it 
places  together  those  elements  which  are  very  closely  related, 
and,  in  addition,  adapts  each  group  to  the  time  of  one  lesson ; 
which  may  be  repeated  if  desirable. 

In  the  Part  devoted  to  gravimetric  and  volumetric  analysis, 
the  examples  are  limited  in  number ;  but  this  much  is  intended 
to  give  the  student  an  opportunity  to  learn  the  methods  with 
the  aid  of  an  instructor,  so  that  he  may  afterwards  pursue  the 
subject  alone,  with  the  aid  of  a  book  like  Fresenius's  Quanti- 
tative Analysis. 

In  conclusion,  it  is  but  just  to  state  that  many  of  the  works 
on  Qualitative  Analysis  have  been  consulted. 

Those  of  Attfield  and  Muter  furnished  many  valuable  hints, 
but  Fresenius's  admirable  work  has  been  used  as  authority, 
and,  to  those  who  wish  to  pursue  the  subject  in  greater  detail 

than  here  offered,  it  is  recommended  for  reference. 

H.  T. 

Philadelphia,  August  igtk,  1883. 


PREFACE  TO  THE  SECOND  EDITION. 


Only  two  important  changes  have  been  thought  advisable 
in  preparing  the  second  edition  of  this  book. 

I.  The  addition  of  equations  explaining  the  more  important 
reactions  which  occur  in  the  action  of  reagents  on  the  various 
salts.     The  typical  ones  are  thus  expressed,  and  the  others 
are  intended  to  be  written  by  the  student  without  aid. 

It  is  believed,  where  there  is  time,  that  the  construction  of 
equations  by  the  student  will  be  of  great  service  to  him,  and 
prevent  too  hasty  performance  of  the  reactions.  It  is  the 
habit  of  the  author  to  question  the  student  closely  on  the 
tests  of  each  element  in  a  group,  before  allowing  him  to 
proceed  to  the  separation.  This  has  been  found  especially 
useful  in  causing  the  student  to  think  for  himself,  and  to 
devise,  from  his  knowledge  of  the  reactions,  a  method  for 
separation  of  the  group. 

II.  While  the  detection  of  acids  in  practice  is  comparatively 
simple,  it  has  been  found  that  it  is  the  source  of  much  diffi- 
culty to  the  beginner.     This  arises  from  the  absence  of  a 
systematic  grouping  of  the  acids,  and  the  separation  of  a  few 
before  attempting  all. 

A  plan  similar  to  that  adopted  for  bases  is  now  proposed, 
which  will  undoubtedly  be  found  an  aid  to  the  student. 
While  it  may  require  a  little  more  time  to  study  the  acids  in 
this  manner,  it  cannot  but  be  of  service  in  facilitating  the 

analysis  of  salts. 

ix 


X  PREFACE   TO    THE   SECOND   EDITION. 

One  word  must  be  said  in  explanation  of  the  use  of 
formulas  and  symbols  for  the  reagents,  instead  of  the  full 
chemical  name.  While  it  may  be  inconvenient  at  first,  the 
student  soon  learns  them,  almost  without  effort,  and  thus 
becomes  familiar  with  what  he  should  know,  but  should  not 
be  compelled  to  memorize. 

The  author's  thanks  are  extended  to  those  who  have  kindly 
furnished  suggestions  for  this  edition,  and  also  to  the  great 
number  who,  without  exception,  favorably  criticised  the  first 
edition. 

The  exhaustion  of  the  first  edition  in  less  than  a  year  is 
satisfactory  evidence  that  an  American  book  on  this  subject, 
adapted  especially  to  the  requirements  of  pharmaceutical  and 
medical  students,  was  needed.  H.  T. 

Philadelphia,  July  jth,  1886. 


CONTENTS. 


PACK 


PART  FIRST— PRACTICAL  CHEMISTRY,      .        .        .17 

SECTION  I— PREPARATION  AND  PROPERTIES  OF  GASES,     .        17 
Hydrogen,  Chlorine,  Hydrochloric  Acid,  Oxygen,  Nitrogen,  Ammo- 
nia, Nitric  Acid,  Carbon  Dioxide, 17~25 

SECTION  II — PREPARATION  OF  SALTS,        ....        26 

Potassium  Chloride,  Potassium  and  Sodium  Tartrate,  Ammonium 
Nitrate,  Ammonium  Oxalate,  Calcium  Phosphate,  Magnesium 
Sulphate,  Magnesium  Carbonate,  Magnesium  Oxide,  Alum- 
inium Hydrate,  Ferrous  Sulphate,  Ferric  Sulphate,  Ferric 
Hydrate,  Copper  Sulphate,  Lead  Acetate,  .  .  .  26-30 

PART  SECOND— QUALITATIVE  ANALYSIS,  .  33 
SECTION  I— BASES, 33 

GROUP  I — Reactions  of  Potassium,  Sodium,  Lithium,  Ammonium,  33 
Summary  of  Group  I,  .  .  .  ,  .  .  .  .  -35 
Analysis  of  Group  I  ........  35 

GROUP  II — Reactions  of  Barium,  Strontium,  Calcium,  Magne- 
sium,  35-37 

Summary  of  Group  II,  ........     3^ 

Analysis  of  Group  II,        ........         38 

Chart  for  Analysis  of  Groups  I  and  II,         .         .         .         .         -39 

Precautions  and  Observations  on  Chart,          ....         39 

GROUP  III — Reactions  of  Manganese,  Zinc,  Cobalt,  Nickel,  .  40-42 
Summary  and  Analysis  of  Group  III,  43 

Chart  for  Analysis  of  Groups  I  to  III,  inclusive,  .         .         -44 

xi 


XI 1  CONTENTS. 

PAGE 

GROUP  IV — Reactions  of  Iron,  Cerium,  Aluminium,  Chromium,   45-47 
Summary  and  Analysis  of  Group  IV,  .....     48 

Chart  for  Analysis  of  Groups  I  to  IV,  inclusive,       ...         49 
Precautions  and  Observations  on  Chart,       .         .         .         .  50 

GROUP  V — Reactions  of  Arsenic,  Antimony,  Tin,  Gold,  Platinum,  50-54 

Summary  and  Analysis  of  Group  V, 55 

Chart  for  Analysis  of  Groups  I  to  V,  inclusive,         ...         56 
Precautions  and  Observations  on  Chart,       .         .         .         .         -57 

GROUP   VI— Reactions    of   Mercury(ic),    Bismuth,    Copper,   Cad- 
mium,                57-59 

Summary  and  Analysis  of  Group  VI, 60 

Chart  for  Analysis  of  Groups  I  to  VI,  inclusive,           .         .         .61 
Precautions  and  Observations  on  Chart, 62 

GROUP  VII— Reactions  of  Silver,  Mercury (ous),  Lead,    .        .  62-64 

Summary  and  Analysis  of  Group  VII, 64 

Chart  for  Analysis  of  Groups  I  to  VII,  inclusive,    ...  -65 

Precautions  and  Observations  on  Chart,           ....  66 

SECTION  II— ACIDS,         .        .        .        .        .        .        .        .67 

GROUP  I — Reactions  of  Acids  Hydrochloric,  Hydrobromic,  Hydri- 

odic,  Hydrofluoric,  Hydrocyanic, 67-69 

Summary  of  Group  I,       ........         69 

Analysis  of  Group  I,     .         .         .         .         .         .         .         .         .69 

GROUP  II— Hypochlorous,  Chloric,  Water,  Hydrates,  Oxides,  Hy- 
drosulphuric,   Sulphurous,   Sulphuric,   Thiosulphuric,    Nitric, 
Hypophosphorous,  Orthophosphoric,  Pyrophosphoric,   Meta- 
phosphoric,  Boric,  Carbonic,  Silicic,         ....        70-75 
Analysis  of  Group  II, 75 

GROUP  III — Acetic,  Valerianic,  Stearic,  Oleic,  Lactic,  Oxalic,  Suc- 
cinic,   Malic,   Tartaric,   Citric,   Carbolic,   Benzoic,   Salicylic, 

Gallic,  Tannic, 77~8o 

Analysis  of  Group  III, 81 


CONTENTS.  Xlll 

PAGE 

SECTION  III— DETECTION  OF  BASES  AND  ACIDS,       .        .  82-86 
Chart  of  Solubilities, 87 

SECTION  IV — REACTIONS  AND  TESTS  OF  ORGANIC  COMPOUNDS,    88 

Alkaloids,    .         . 9I-93 

Neutral  Principles, 94 

PART  THIRD— QUANTITATIVE  ANALYSIS,      .        .    97 
SECTION  I— GRAVIMETRIC  ESTIMATION,     ....         97 

Preliminary  Directions, 97 

Examples  in  Gravimetric  Estimation,  Barium,  Chlorine,  Copper, 
Sulphuric  Acid,  Potassium,  Nitric  Acid,  Calcium,  Carbonic 
Acid, 98-101 

SECTION  II— VOLUMETRIC  ESTIMATION,         .        .        .        .  102 

Examples,  Oxalic  Acid,  Sodium  Hydrate,  Potassium  Bichromate, 

Iodine,  Sodium  Hyposulphite,  Silver  Nitrate,  .  102-106 
Table  of  Elements,  Symbols  and  Atomic  Weights,  .  .  .  107 
Index, 109 


PART  FIRST. 


PRACTICAL 


fiti 


PRACTICAL  AND  ANALYTICAL 

CHEMISTRY. 


PART  I.— PRACTICAL. 

SECTION  I. 

PREPARATION  AND  PROPERTIES  OF  GASES. 


HYDROGEN,  H. 

Preparation. — Place  a  few  fragments  of  zinc  in  a  flat-bottom 
flask  of  about  one-fourth  liter  capacity ;  cover  the  zinc  with 
water  and  adapt  a  cork,  through  which  pass  two  tubes  (Fig.  i), 
one  just  reaching  through  the  cork  and  bent  so  the  long  end 
may  be  dipped  under  water,  the  other  running  directly  from 
a  short  distance  above  the  cork  nearly  to  the  bottom  of  the 
flask,  so  as  to  be  below  the  surface  of  the  liquid.  The  upper 
end  of  this  tube  should  have  a  small  funnel  placed  in  it,  or  be 
enlarged  by  softening  in  the  flame,  inserting  and  revolving  a 
file  or  similar  instrument,  previously  warmed. 

Add,  slowly,  a  small  quantity  of  sulphuric  acid,  through 
this  tube,  and  notice  an  immediate  effervescence,  with  the 
escape  of  bubbles  through  the  water  in  which  the  exit  tube 
dips.  Fill  a  test  tube  with  water,  and,  keeping  the  open  end 
under  the  liquid,  bring  it  over  the  tube,  so  as  to  collect  the 
gas.  When  full,  close  with  the  thumb,  and,  bringing  the  mouth 
of  the  tube  near  a  flame,  quickly  remove  the  thumb  and  allow 
the  gas  to  ignite.  It  will  burn  quietly  if  the  gas  be  pure,  but 
with  a  slight  explosion  if  it  be  mixed  with  air. 

Properties. — This  gas  is  Hydrogen,  and  its  physical  prop- 
erties may  now  be  studied  by  observing  that  it  is  insoluble  in 
water,  and  without  odor,  color  or  taste. 
2  17 


18 


PRACTICAL   CHEMISTRY. 


EXPERIMENT  I.  Collect  a  tube  full,  and  holding  it,  covered, 
in  a  vertical  position,  bring  a  lighted  taper  a  short  distance 
above  its  mouth  and  remove  the  cover ;  the  gas  will  ignite, 
showing  its  great  levity. 

EXPERIMENT  II.  Another  tube,  similarly  filled,  is  held  in 
an  inverted  position,  and  the  cover  removed ;  it  will  be  found, 
even  after  the  lapse  of  some  time,  that  the  hydrogen  at  the 
mouth  of  the  tube  may  be  ignited,  thus  demonstrating  that  the 
gas  is  too  light  to  come  down  and  out  the  mouth  of  the  tube. 
These  two  experiments  have  also  demonstrated  the  combus 

FIG.  i. 


tibility  of  the  gas,  which,  when  pure,  burns  quietly,  with  a 
colorless  flame.  If,  however,  it  be  mixed  with  air  and  flame 
applied,  a  violent  explosion  ensues.  Therefore,  the  tube  from 
the  generator  should  never  be  brought  near  a  flame  until  it  is 
certain  all  the  air  has  been  expelled.  This  is  determined  by 
trying  a  test  tube  full;  if  it  burn  quietly,  the  jet  may  be  lighted. 
This  precaution  should  always  be  observed. 

EXPERIMENT  III.  On  bringing  a  lighted  taper  to  the  mouth 
of  a  tube  full  of  hydrogen,  the  gas  is  ignited,  but  on  pushing 
the  burning  taper  up  into  the  gas  its  flame  is  extinguished, 


PREPARATION    AND    PROPERTIES    OF    GASES. 


19 


thus  showing  that  while  hydrogen  is  combustible  it  is  not  a 
supporter  of  combustion. 

EXPERIMENT  IV.  Continue  the  addition  of  acid  to  the 
zinc  until  the  latter  is  nearly  all  dissolved;  disconnect  the 
apparatus,  pour  the  liquid  on  a  filter,  collect  the  filtrate  in 
a  small  beaker  or  evaporating  dish,  concentrate  and  set  aside 
for  twenty-four  hours,  to  crystallize.  These  crystals  are  zinc 

FIG.  2. 


sulphate,  ZnSO4,  the  result  of  a  combination  of  the  sulphuric 
acid  and  the  zinc,  as  follows  : — 

Zn  -f  H2SO4  =  ZnSO4  +  H2. 

CHLORINE,  Cl. 

Preparation. — In  a  flask,  arranged   so   that  heat  may  be 
applied  (Fig.  2),  place  a  small  quantity  of  manganese  dioxide, 


20  PRACTICAL   CHEMISTRY. 

add  hydrochloric  acid,  agitate  well,  to  moisten  all  the  powder 
on  the  bottom,  and  apply  heat.  A  yellowish-green  gas  is 
evolved,  which,  being  somewhat  soluble  in  water,  may  be 
collected  by  downward  displacement,  that  is,  by  running  the 
delivery  tube  to  the  bottom  of  the  receptacle,  loosely  covered, 
the  heavy  gas  displaces  the  lighter  air. 

If  the  evolution  be  moderately  active  it  may  be  collected 
over  warm  water,  as  only  a  small  loss  occurs.  Care  should 
be  taken  to  avoid  inhaling  the  gas,  by  passing  it,  when  not 
collecting,  into  a  solution  of  potassium  or  sodium  hydrate. 
The  following  expresses  the  reaction  in  symbols  : — 

MnO2  +  4HC1  =  MnCl2  +  2H2O  +  C12. 

Properties. — EXPERIMENT  I.  Pass  the  gas  into  water;  it 
is  absorbed;  if  this  be  continued  until  the  water  is  saturated,  it 
will  be  found  to  have  absorbed  about  twice  its  volume  of  the  gas  ; 
the  Aqua  Chlori  of  the  Pharmacopoeia  is  the  resulting  product. 

EXPERIMENT  II.  A  tube  full  of  the  gas  held  with  mouth 
upward,  and  a  lighted  taper  applied,  fails  to  ignite.  Push  the 
taper  into  the  gas ;  it  is  extinguished,  or  only  burns  with  a 
small,  dense,  smoky  flame,  the  result  of  a  combination  of  the 
chlorine  with  the  hydrogen  of  the  wax,  liberating  the  carbon. 

EXPERIMENT  III.  Into  a  tube  full  of  the  gas  put  a  piece 
of  brightly  dyed  calico,  previously  moistened ;  it  is  rapidly 
bleached.  Writing  on  paper  is  similarly  decolorized,  but 
printing  is  not  affected,  as  it  contains  carbon,  in  the  form  of 
lampblack,  which  is  not  acted  on  by  the  gas.  If  the  experi- 
ments be  made  with  chlorine  which  has  been  passed  through 
sulphuric  acid  to  dry  it,  and  the  materials  are  not  moistened, 
no  decoloration  takes  place. 

The  process  of  bleaching  by  chlorine  is  one  of  oxidation  ;  it 
combines  with  the  hydrogen  of  the  water,  forming  hydro- 
chloric acid,  while  the  liberated  oxygen  in  the  nascent  state 
readily  attacks  the  coloring  matter,  water  and  a  colorless  com- 
pound resulting. 

HYDROCHLORIC   ACID,  HC1. 

Preparation. — Collect  one  test  tube  full  of  hydrogen  and 
one  of  chlorine,  bring  their  mouths  together  (the  hydrogen 
tube  above  with  mouth  down,  as  it  is  lighter),  turn  over  once 


PREPARATION  AND  PROPERTIES  OF  GASES.         21 

or  twice,  so  as  to  thoroughly  mix,  and  open  their  mouths  to 
a  flame;  a  sharp  report  will  occur,  with  the  development  of 
strongly  acid  fumes,  which  will  be  recognized  by  future  tests 

as  hydrochloric  acid : — 

H2  +  C12  =  2HC1. 

To  prepare  a  quantity  of  it,  the  apparatus  used  for  the 
preparation  of  chlorine  serves  best.  Put  into  the  flask  some 
sodium  chloride  (common  salt)  add  sulphuric  acid  slowly,  and 
when  the  evolution  of  gas  ceases  apply  a  gentle  heat.  Collect 
by  downward  displacement  or  over  mercury : — 
NaCl  +  H2SO4  r=  NaHSO4  -f  HC1. 

The  above  expresses  the  reaction  when  an  excess  of  acid 
has  been  used,  which  is  preferable,  as  the  resulting  acid  sodium 
sulphate  is  easily  dissolved  out  of  the  flask  with  water.  On 
the  large  scale  the  following  more  economical  method  is 
used  : — 

(NaCl)2  +  H2SQ4  =  Na2S04  +  (HC1)2. 

Properties. — The  pungent,  suffocating  odor  and  freedom 
from  color  are  noted  with  its  production. 

EXPERIMENT  I.  A  test  tube  or  jar  of  the  gas  placed  with 
the  open  mouth  under  water  will  so  rapidly  dissolve  that  the 
liquid  rises  in  the  vessel.  This  solution  of  the  gas  in  water, 
when  of  the  proper  strength,  is  the  Acidum  Hydrochloricum, 
U.  S.  P. 

EXPERIMENT  II.  A  piece  of  moistened  blue  litmus  paper 
held  in  a  tube  of  the  gas  is  instantly  reddened. 

EXPERIMENT  III.  A  lighted  taper  applied  to  the  gas  fails 
to  ignite  it,  and  is  extinguished  if  lowered  into  it. 

EXPERIMENT  IV.  Bring  a  rod  moistened  with  ammonia 
over  the  mouth  of  a  tube  full  of  the  gas :  dense  white  fumes 
of  ammonium  chloride  are  formed. 

The  mixture  remaining  in  the  flask  dissolved  in  warm 
water,  treated  with  sodium  carbonate  so  long  as  effervescence 
occurs,  to  neutralize  the  excess  of  sulphuric  acid,  concen- 
trated, filtered  and  set  aside  to  crystallize,  yields  sodium  sul- 
phate (Glauber  salt). 

OXYGEN,  O. 

Preparation. — Place  a  few  crystals  of  potassium  chlorate 
in  a  test  tube,  adapt  a  delivery  tube  long  enough  to  reach 


22  PRACTICAL    CHEMISTRY. 

under  the  surface  of  some  water  near  by.  Apply  a  steady 
flame ;  as  soon  as  the  bubbles  of  gas  escape  freely  and  the 
air  has  been  expelled,  bring  a  test  tube  filled  with  water  over 
the  escaping  gas,  and  collect.  It  is  Oxygen,  produced  from 
the  potassium  chlorate  by  heat,  according  to  the  following 

reaction : — 

KC103  =  KC1  +  08. 

Properties. — The  appearance  and  insolubility  of  the  gas  in 
water  are  noted  as  it  is  collected. 

EXPERIMENT  I.  The  gas  is  not  ignited  by  the  application 
of  a  lighted  taper.  The  taper,  however,  will  burn  with  greatly- 
increased  energy  if  it  be  plunged  into  the  gas.  If  the  flame 
be  extinguished  and  the  taper  again  brought  into  the  gas, 
provided  a  spark  remain,  the  taper  is  rekindled. 

EXPERIMENT  II.  A  piece  of  charcoal,  previously  ignited,  is 
lowered  into  the  gas ;  a  rapid  combustion  ensues,  and  the 
charcoal  disappears.  Pour  some  lime  water  into  the  tube, 
agitate  well ;  a  white  precipitate  of  calcium  carbonate  is  pro- 
duced. If  this  be  tried  with  oxygen  previous  to  the  burning 
of  the  charcoal,  no  precipitate  will  be  formed.  A  number  of 
other  substances,  as  sulphur,  phosphorus,  and  even  iron,  when 
once  kindled,  will  burn  in  oxygen  with  great  brilliancy,  form- 
ing characteristic  oxides. 

EXPERIMENT  III.  Take  two  test  tubes,  one  about  twice  the 
capacity  of  the  other.  Fill  the  larger  with  hydrogen  and  the 
smaller  with  oxygen,  bring  their  mouths  together,  and,  after 
turning  once  or  twice,  to  thoroughly  mix  their  contents,  open 
them  and  apply  flame.  A  sharp  report  is  caused  by  the  com- 
bination to  form  water. 

If  more  than  two  volumes  of  hydrogen  to  one  of  oxygen 
are  present,  the  surplus  remains  uncombined;  if  oxygen  is 
present  in  greater  proportion,  the  excess  of  it  remains. 

NITROGEN,  N. 

Preparation. — The  usual  method  of  preparing  this  gas  is 
to  deprive  air  of  its  oxygen,  leaving  the  nitrogen  pure.  This 
is  best  accomplished  by  placing  a  small  fragment  of  phos- 
phorus on  a  cork  covered  with  some  fireproof  material.  Float 
the  cork  and  phosphorus  on  water,  ignite  the  latter  and  bring 


PREPARATION    AND    PROPERTIES    OF    GASES.  23 

over  it  a  bell-jar.  The  phosphorus  combines  with  the  oxygen, 
converting  it  into  phosphorus  pentoxide,  P2O5,  which  dissolves 
in  the  water  present,  thus  leaving  the  nitrogen  pure.  A  better 
method  for  obtaining  larger  quantities  is  to  heat  a  mixture  of 
potassium  nitrite  and  ammonium  chloride  dissolved  in  water. 
When  the  reaction  begins,  the  temperature  must  be  carefully 
watched,  in  order  to  prevent  the  too  rapid  evolution  of  the 

gas:— 

KNOa  +  NH4C1  =  KC1  +  2  H2O  +  N2. 

Properties. — The  physical  properties  have  been  observed 
during  its  preparation  and  collection.  In  regard  to  chemical 
properties,  it  is  inert  in  the  free  state.  Its  compounds,  how- 
ever, are  very  energetic. 

AMMONIA,  MH3. 

Preparation. — In  a  test  tube  or  evaporating  dish  mix  equal 
quantities  of  powdered  calcium  oxide  (quicklime)  and  ammo- 
nium chloride,  with  a  few  drops  of  water;  the  odor  of  ammonia 
will  be  immediately  developed : — 

CaO  +  (NH4C1)2  =  CaCl2  -f  H2O  ±  (NH8)2. 

In  smaller  quantities  the  gas  may  be  recognized  by  holding 
over  the  mixture  a  strip  of  moistened  red  litmus  paper ;  it  will 
slowly  become  blue ;  or  similarly  hold  a  glass  rod  moistened 
with  hydrochloric  acid ;  dense  white  fumes  of  ammonium 
chloride  will  form.  To  prepare  larger  quantities  of  the  gas, 
heat  the  ordinary  water  of  ammonia,  which,  at  a  comparatively 
low  temperature,  gives  it  off  freely.  If  it  be  desired  to  dry 
the  gas,  it  must  be  passed  over  quicklime.  Ammonia  is  col- 
lected by  upward  displacement,  that  is,  by  passing  the  delivery- 
tube  upward  into  a  jar  or  test  tube  inverted  over  it;  being 
lighter  than  air  the  latter  is  diplaced. 

Properties. — EXPERIMENT  I.  Place  a  vessel  filled  with 
ammonia  gas,  mouth  downward,  into  some  water,  and  agitate 
slightly;  the  water  will  rise  in  the  vessel  rapidly,  nearly  filling 
it,  showing  the  great  solubility  of  the  gas  in  water.  The  other 
physical  properties,  as  color,  odor,  etc.,  have  been  noted  during 
its  collection. 

EXPERIMENT  II.  On  applying  a  lighted  taper  to  the  gas  it 
does  not  burn;  if,  however,  it  be  mixed  with  oxygen  it  will 


24  PRACTICAL    CHEMISTRY. 

ignite  readily.     On  introducing  the  taper  into  the  gas  it  is 
extinguished. 

NITRIC   ACID,  HNO3. 

Preparation. — Place  a  small  quantity  of  potassium  nitrate 
in  a  test  tube,  and  cover  it  with  strong  sulphuric  acid.  Apply 
a  gentle  heat ;  brown,  strongly  acid  fumes  are  given  off. 
Dilute  with  a  little  water  and  add  indigo  solution;  it  is  decolor- 
ized. This  is  a  characteristic  test  for  nitric  acid. 

To  prepare  a  larger  quantity  a  retort  is  used,  to  which  is 
adapted  a  glass  receiver.  The  nitric  acid  distills  over  on  the 
application  of  a  moderate  heat,  forming  a  reddish-yellow 
liquid,  which  rapidly  attacks  and  destroys  organic  matter. 

Two  reactions  may  be  employed  to  represent  the  produc- 
tion of  nitric  acid,  depending  on  the  relative  quantity  of  the 
materials  used.  In  the  first  case  an  excess  of  sulphuric  acid 

gives — 

KNO3  -f  H2SO4  =  KHSO4  +  HNO3. 

In  the  second  case  just  a  sufficient  amount  of  sulphuric  acid 
is  used  to  decompose  the  potassium  nitrate — 

(KNO3)2  +  H2SO4  =  K2SO4  +  (HNO3)2. 

On  the  large  scale  sodium  nitrate  is  now  employed  in  place  of 
the  potassium  salt,  on  account  of  its  lower  price. 

CARBON    DIOXIDE,  CO2. 

Preparation. — The  flask  used  in  the  preparation  of  hydro- 
gen will  serve  for  making  carbon  dioxide.  A  few  pieces  of 
marble  are  placed  in  the  flask,  covered  with  water,  and  hydro- 
chloric acid  added.  A  brisk  effervescence  ensues,  and  the  gas 
being  somewhat  soluble  in  water  is  collected  by  downward 
displacement — 

CaC03-+  (HC1)2  =  CaCl2  +  H2O  +  CO2. 

Sulphuric  acid  should  not  be  used,  as  it  forms  an  insoluble 
calcium  sulphate  which  is  very  difficult  to  remove  from  the  flask. 

Properties. — EXPERIMENT  I.  Pour  some  clear  lime  water 
into  a  jar  of  the  gas  and  agitate ;  the  solution  immediately 
becomes  cloudy,  owing  to  formation  of  insoluble  calcium  car- 
bonate. If  more  gas  be  passed  into  the  mixture  it  will  become 
clear  again,  on  account  of  the  solubility  of  the  precipitate  in 
carbonic  acid. 


PREPARATION  AND  PROPERTIES  OF  GASES.         25 

EXPERIMENT  II.  Add  solution  of  potassium  hydrate  to  a 
jar  of  the  gas,  close  and  shake  well ;  the  gas  is  absorbed  by 
the  alkali,  as  may  be  shown  by  placing  the  mouth  of  the  jar 
under  water  and  removing  the  stopper,  when  the  water  will 
rush  in,  nearly  filling  it. 

EXPERIMENT  III.  A  lighted  taper  lowered  into  the  gas  is 
immediately  extinguished.  The  same  result  is  accomplished 
by  opening  the  vessel  some  distance  above  the  flame  and 
allowing  the  gas  to  flow  down  upon  it.  This  latter  experi- 
ment also  illustrates  the  great  density  of  the  gas,  which  is 
twenty-two  times  heavier  than  hydrogen. 


26  PRACTICAL    CHEMISTRY. 

SECTION  II. 

PREPARATION    OF    SALTS. 

POTASSIUM   CHLORIDE,  KC1. 

Preparation. — One  or  two  cubic  centimeters  of  hydro- 
chloric acid,  diluted  with  three  or  four  times  its  bulk  of  water, 
are  placed  in  a  small  beaker  glass,  and  potassium  carbonate 
added  so  long  as  effervescence  occurs,  and  until  after  boiling 
(to  remove  CO2)  the  solution  is  neutral  to  litmus  paper,  that 
is,  when  the  blue  litmus  paper  is  not  changed  to  red  nor  the 
red  changed  to  blue.  Evaporate  to  a  small  bulk  and  set  aside 
to  crystallize.  The  cubic  crystals  which  separate  after  standing 
twenty-four  hours  may  be  collected  on  filter  paper  and  dried 
at  a  moderate  temperature  : — 

K2C03  +  (HC1)2  =  (KC1)2  +  H20  +  C02. 

Potassium  chloride  is  rarely  prepared  in  this  manner,  except  for  practice,  as  it 
occurs  largely  in  nature,  and  is  used  for  preparing  many  other  potassium  salts. 

POTASSIUM   AND   SODIUM   TARTRATE, 

KNaC4H4O64H2O. 

(ROCHELLE  SALT.) 

Preparation. — Heat,  in  a  porcelain  capsule,  a  solution  of 
sodium  carbonate,  and  add  to  it  potassium  bitartrate  until 
effervescence  ceases,  and  the  solution  (after  the  escape  of  CO2) 
is  neutral  to  litmus.  On  filtering  and  cooling,  crystals  of 
Rochelle  salt  are  deposited,  rapidly  and  in  small  crystals  if  the 
solution  be  concentrated,  but  slowly  and  in  much  larger  ones 
if  the  solution  be  dilute : — 

Na2C03  +  (KHC4H406)2  i=  (KNaC4H4O6)2  -f  H26  +  CO2. 

Note  on  Calculation. — In  order  to  calculate  the  amount  of  each  salt  to  use  in 
the  above  process,  we  notice  the  number  of  molecules  of  each  employed,  and 
multiply  this  by  the  sum  of  the  atomic  weights. (that  is,  by  the  molecular  weights). 
In  the  above  case  one  molecule  of  Na2CO3  =  106,  and  two  molecules  of 
KHC4H4O6  =  2  X  1 88. 1  =  376.2.  Therefore  every  106  parts  of  anhydrous 
sodium  carbonate  require  376.2  parts  of  potassium  bitartrate,  to  form  Rochelle 
salt.  If  we  have  50  grams  of  anhydrous  sodium  carbonate  and  wish  to  convert 
it  into  Rochelle  salt,  we  use  the  following  formula :  As  106  :  376.2  :  :  50  grams  : 
number  of  grams  of  potassium  bitartrate  required,  =  177.4  grams. 


PREPARATION   OF   SALTS.  27 

AMMONIUM    NITRATE,  NH4NO3. 

Preparation. — Add  to  about  20  c.c.  of  dilute  nitric  acid,  in 

a  beaker  glass,  sufficient  ammonia  water  to  give  it  a  distinct 

ammoniacal  odor;    filter,  concentrate,  keeping  the  ammonia 

in  slight  excess,  and  set  aside  in  cool  place   for  crystals   to 

form  : — 

NH4OH  -f  HNO3  =  NH4NO3  -f  H2O. 

Properties. — These  crystals  contain  twelve  molecules  of 
water  of  crystallization,  which  it  is  desirable  to  get  rid  of  be- 
fore using  the  salt.  By  exposure  to  a  temperature  of  155°  C. 
the  water  gradually  escapes,  and  the  fused  or  granulated  salt 
is  ready  to  be  converted  into  nitrogen  monoxide  (laughing  gas), 
which  takes  place  at  about  185°  C.  according  to  the  following 

reaction : — 

NH4N03=N20  +  (H20)2. 

AMMONIUM    OXALATE,  (NH4)2C2O4. 

Preparation. — Dilute  20  c.c.  of  solution  of  ammonia  with 
twice  its  bulk  of  water,  add  a  solution  of  oxalic  acid  until 
neutral,  concentrate  slightly,  filter  and  set  aside  to  crystallize. 
The  crystals  may  be  collected  on  a  filter,  and  another  crop 
obtained  by  concentrating  the  "mother  liquor  ": — 
(NH4OH)2  +  H2C204  ==  (NH4)2C204  +  (H2O)2. 

Ammonium  carbonate  is  sometimes  used  for  combining  with  the  oxalic  acid, 
but  the  neutralization  is  not  as  easily  effected,  besides  it  is  not  desirable  on  the 
ground  of  economy. 

CALCIUM  PHOSPHATE,  Ca3(PO4)2. 
Preparation. — Finely  powdered  bone  ash  is  digested  for 
a  short  time  with  'diluted  hydrochloric  acid.  The  solution 
filtered,  boiled,  filtered  again,  if  necessary.  The  filtrate  is 
treated  with  ammonia  until  it  smells  strongly  of  it.  Collect 
the  precipitate  on  a  filter,  wash  by  pouring  on  warm  water 
until  the  washings  are  tasteless,  and  dry  at  a  low  temperature. 
The  resulting  powder  is  calcium  phosphate,  which  exists  in 
the  bone  ash  and  is  dissolved  by  hydrochloric  acid,  forming 
acid  calcium  phosphate,  as  follows  : — 

Ca3(P04)2  +  (HC1)4  =  CaH4(P04)2  +  (CaCl2)2. 

From  this  solution  it  is  precipitated  by  ammonia,  as  follows : — 
CaH4  (P04)2  +  (CaCl2)2  -f  (NH4OH)4  =  Ca3(PO4)2  +  (NH4C1)4. 


OF   THTC 

UNIVERSITY  H 


28 


PRACTICAL    CHEMISTRY. 


In  addition  to  the  ordinary  apparatus  with  which  a  student 
supplies  himself,  there  is  required  a  wash  bottle  (Fig.  3),  which 
it  is  well  for  every  student  to  construct  for  himself,  as  it 
furnishes  him  valuable  practice  in  cutting  and  bending  glass 
tubing.  This  bottle  is  used  in  washing  all  precipitates,  and  is 
convenient  as  a  water  supply,  which  may  be  kept  hot,  if  desired. 

MAGNESIUM   SULPHATE,  MgSO4.;H2O. 
Preparation. — To  about  5   c.c.  of  sulphuric  acid,  diluted 
with  five  or  six  times  its  volume  of  water,  heated  in  a  capsule, 

FIG.  3. 


add  powdered  magnesium  carbonate  until  effervescence  ceases, 
and  filter.     Concentrate  and  set  aside  to  crystallize. 

(MgC03)4Mg(OH)2.5H20  +  (H2S04)5  =  (MgSO4)5  +  (H.O)^ 


MAGNESIUM    CARBONATE,  (MgCO3)4Mg(OH)2.5H2O. 

Preparation.  —  On  mixing  solutions  of  magnesium  sulphate 
and  sodium  carbonate  and  boiling,  we  get  magnesium  car- 
bonate precipitated,  while  carbon  dioxide  escapes.  The  pre- 
cipitate is  very  variable  in  its  composition,  depending  on  the 
concentration  of  the  solutions.  When  the  U.  S.  P.  product  is 
obtained  the  following  equation  expresses  the  reaction  :  — 

(MgS04)5  +  (Na2C03)5  +  (H20)6  = 
(MgC03)4Mg(OHJ2.5H20  +  (Na2S04)5  +  CO2. 


PREPARATION    OF    SALTS.  29 

The  precipitate  washed  with  hot  water  and  dried,  serves  for 
the  following  example  of  a  compound  prepared  by  ignition. 

MAGNESIUM    OXIDE,  MgO. 

Preparation. — Heat  some  of  the  magnesium  carbonate, 
prepared  in  the  above  reaction,  in  a  porcelain  crucible  until, 
on  taking  out  a  small  portion,  placing  in  a  test  tube  with  a 
little  water,  heating  to  remove  air  bubbles,  and  adding  a  drop 
or  two  of  hydrochloric  acid,  no  effervescence  is  produced. 
This  will  require  some  time,  and  great  care  is  necessary  to 
determine  when  the  powder  fails  to  give  an  effervescence  with 
the  acid. 

(MgC03)4Mg(OH)25H20  =  (MgO)5  +  (CO2)4  +  (H2O)6. 
Zinc  oxide  may  be  prepared  in  a  similar  manner,  from  zinc  carbonate.     This 
differs  from  the  magnesium  oxide  by  being  yellow  while  hot,  and  very  pale  yellow 
'when  cold. 

ALUMINIUM  HYDRATE,  A12(OH)6. 
Preparation. — To  a  solution  of  alum  add  a  solution  of 
sodium  carbonate  and  boil.  Allow  the  precipitate  to  settle, 
decant  the  clear  supernatant  liquid  on  a  filter,  add  more  hot 
water  to  the  precipitate  and  again  decant.  Collect  the  pre- 
cipitate on  the  filter,  wash  well  with  hot  water  and  dry  ;  the 
resulting  white  powder  is  Aluminii  Hydras,  U.  S.  P. 

A12(S04)3,  K2S04  -f  (Na2C03)3  +  (H2O)3  = 
A12(OH)6  +  K2S04  +  (Na2S04)3  +  (CO2)3. 

FERROUS   SULPHATE,  FeSO4.;H2O. 

Preparation. — Add  enough  dilute  sulphuric  acid  to  some 

iron  filings,  or  wire,  in  a  beaker,  to  cover  them.     AlldWthe 

reaction  to  proceed,  assisted  by  a  little  heat,  until  effervescence 

ceases.     Filter  from  the  excess  of  iron,  concentrate,  filter  and 

crystallize. 

(Fe)2  +  (H2S04)2  =  (FeS04)2  +  (Ha)a. 

These  crystals  should  be  rapidly  dried  and  preserved  in  well  stopped  bottles,  as 
they  quickly  become  converted  into  ferric  sulphate  on  exposure  to  air. 

FERRIC   SULPHATE,   Fe2(SO4)3. 

Preparation. — To  a  strong  solution  of  ferrous  sulphate  add 
one-fourth  its  bulk  of  sulphuric  acid,  heat  to  the  boiling  point 
and  drop  in  nitric  acid  as  long  as  effervescence  is  produced 


30  PRACTICAL    CHEMISTRY. 

and    until   the    resulting  liquid  becomes  of  a  clear  reddish- 
brown  color. 

(FeS04)6  +  (H2S04)3  +  (HN03)2  =  (Fe2(SO4)3)3  +  N2O2  -f  (H2O)4. 
This  is  the  Liquor  Ferri  Tersulphatis  of  the  Pharmacopoeia ;  and  is  the  most 
convenient  compound  to  use  in  the  preparation  of  some  of  the  other  iron  salts. 

FERRIC    HYDRATE,  Fe2(OH)6. 

Preparation. — Dilute  some  of  the  above  ferric  sulphate 
solution  with  an  equal  bulk  of  water,  add  solution  of  am- 
monia until,  after  stirring,  it  smells  strongly.  The  resulting 
precipitate  is  ferric  hydrate,  the  well-known  antidote  to  arsenic. 
When  needed  for  this  purpose,  it  is  sufficient  to  pour  the 
mixture  on  a  muslin  strainer,  wash  once  or  twice  until  the 
saline  taste  nearly  disappears  from  the  washings,  when  the 
compound  is  ready  for  use. 

Fe2(S04)3  +  (NH4OH)6  =  Fe2(OH)6  +  (  (NH4)2SO4)3.  7 

This  preparation  should  always  be  freshly  prepared  when  wanted  for  use  as  an 
antidote,  as  it  loses  H2O  on  keeping,  becoming  a  mixture  of  ferric  oxide  Fe2O3 
and  hydrate.  This  change  takes  place,  although  more  slowly,  when  the  com- 
pound is  kept  under  water. 

COPPER   SULPHATE,  CuSO4.sH2O. 
Preparation. — Heat  copper  turnings  for  some  time,  with 
strong  sulphuric  acid,  in   a   fume   closet,   until    the   reaction 
ceases.     Dilute  with  water,  filter  and  crystallize. 

Cu  -f  (H2SO4)2  =  CuSO4  +  SO2  +  (H2O)2. 

LEAD   ACETATE,  Pb(C2H8O2)2.3H2O. 

Preparation. — Lead  oxide  (litharge)  is  boiled  with  three  or 
four  times  its  weight  of  acetic  acid,  in  a  capsule,  adding  water 
from  time  to  time,  with  more  acid,  if  necessary,  until  most  of 
the  oxide  has  disappeared.  Filter,  concentrate,  keeping  the 
solution  acid,  and  set  aside  to  crystallize  : — 

PbO  -f  (HC2H3O2)2  =  Pb(C2H3O2)2  +  H2O. 

The  solution  or  crystals  should  not  be  exposed  to  the  fumes  of  the  laboratory, 
for  if  there  be  only  a  small  quantity  of  hydrogen  sulphide  in  the  room,  they  will 
become  black. 


PART  SECOND. 


QUALITATIVE  ANALYSIS. 


PART  II.     QUALITATIVE  ANALYSIS. 

SECTION  I. 

BASES. 


GROUP  I.— POTASSIUM,  SODIUM,  LITHIUM, 
AMMONIUM. 

REACTIONS   OF  POTASSIUM   (K). 
Use  a  solution  of  potassium  chloride  (KC1). 

1.  PtCl4  causes  a  yellow  crystalline  precipitate  of  K2PtCl6, 
soluble  in  excess  of  water. 

PtCl4  +  2KC1  =  K2PtCl6. 

The  delicacy  of  this  reaction  is  increased  by  the  addition  of 
alcohol. 

2.  H2C4H4O6,  in  concentrated  solution,  produces  a  white  crys- 
talline precipitate  of  potassium  acid  tartrate — KHQH4O6, 
soluble  in  excess  of  water,  readily  in  hot  water,  acids  or  potas- 
sium hydrate. 

H2C4H4O6  +  KC1  =  KHC4H4O6  +  HC1. 

The  addition  of  alcohol  and  violent  agitation  facilitate  the 
formation  of  this  precipitate. 

3.  A  fragment  of  potassium  salt  on  the  loop  of  a  platinum 
wire,  held  in  the  colorless  flame  of  a  Bunsen  gas  lamp  imparts 
a  violet  color.     This  reaction  is  interfered  with  by  the  presence 
of  sodium  salts,  which  color  ftie  flame  yellow  ;  organic  matter 
also  colors  the  flame  violet,  and  should  be  removed  by  ignition 
before  testing  for  potassium.     The  yellow  rays  of  sodium  may 
be  destroyed  by  viewing  the  flame  through  blue  glass. 

4.  Potassium  salts  are  not  volatile  at  a  low  red  heat ;  at  a 
white  heat  they  are  slowly  volatilized. 

REACTIONS   OF   SODIUM   (Na). 
Use  a  solution  of  sodium  chloride  (NaCl). 
I.  Sodium   salts  color  the  gas  flame  yellow;  so  delicate  is 
this  reaction  that  the  merest  traces  are  revealed  by  it. 
3  33 


34  ANALYTICAL   CHEMISTRY. 

2.  The  salts  of  sodium  are  not  volatile  at  a  low  red  heat, 
but  slowly  volatilize  at  a  white  heat. 

REACTIONS   OF   LITHIUM  (Li). 
Use  a  solution  of  lithium  chloride  (LiCl). 

1.  Na2HPO4   added    to    a    strong    solution    produces,    on 
boiling,  a  white  precipitate  of  lithium  phosphate — Li3PO4. 

Na2HPO4  -f  3L1C1  =  Li3PO4  +  2NaCl  -f  HC1. 

This  reaction  takes  place  more  readily  when  the  solution  is 
first  made  alkaline  with  NH4OH. 

2.  Lithium  salts  impart  an  intense  crimson  color  to  the  gas 
flame.     This  is  somewhat  interfered  with  by  sodium  salts,  but 
the  yellow  color  of  sodium  may  be  excluded  by  blue  glass, 
which  if  not  too  dark  will  allow  the  crimson  rays  of  lithium 
to   pass   through.      These   must   not   be   confused   with   the 
violet  potassium  rays,  which  will  pass  through  a  deep  blue 
glass. 

3.  Lithium  salts  do  not  volatilize  at  a  low  red  heat,  but  are 
slowly  volatilized  at  a  white  heat. 

REACTIONS.  OF   AMMONIUM  (NH4). 
Use  a  solution  of  ammonium  chloride  (NH4C1). 

1.  PtCl4  produces,  in  strong  solution,  a  yellow  crystalline 
precipitate  of  ammonium  platino-chloride — (NH4)2PtCl6. 

2.  NaOH   on    heating  causes  the  evolution  of  ammonia — 
NH3;   detected   by  the    odor;  by   holding  near  a  glass   rod 
moistened  with  HC1,  which  will  produce  dense  white  fumes 
of  NH4C1 ;  or  by  holding  in  the  mouth  of  the  tube  a  strip  of 
moistened  red  litmus  paper,  when  it  will  immediately  become 
blue.     Care  must  be  taken,  in  this  last  test,  to  prevent  any  of 
the  alkaline  liquid  coming  in  contact  with  the  paper,  as  it 
would  likewise  cause  the  blue  color. 

3.  H2C4H4O6,  added  to  a  concentrated  solution,  produces  a 
white  precipitate  of  ammonium  acid  tartrate — NH4HC4H4O6, 
soluble  in  slight  excess  of  water. 

4.  Ammonium  salts  are  volatile  at  a  low  red  heat. 


BASES.  35 

SUMMARY  OF  TESTS   WITH   SOLUBLE  SALTS   OF  GROUP   I. 


K 

Na 

Li 

NH4 

PtCl4 

Yellow  Precipitate 

No  Precipitate 

No  Precipitate 

Yellow  Precipitate 

H2C4H406 

White  Precipitate 

No  Precipitate 

No  Precipitate 

White  Precipitate 

Flame 

Violet 

Yellow 

Crimson 

None 

Volatility 

Not  Volatile 

Not  Volatile 

Not  Volatile 

Volatile 

DIRECTIONS    FOR   THE   DETECTION   OF   THE 

BASES   IN   A   SOLUTION    CONTAINING 

SOLUBLE   SALTS   OF   GROUP   I. 

To  a  small  portion  of  the  solution  add  NaOH  and  heat : 
NH3  will  be  given  off  if  ammonium  salts  are  present,  and 
may  be  detected  by  the  odor,  or  by  moistened  red  litmus 
paper. 

Evaporate  another  portion  of  the  solution  to  dryness, 
transfer  to  a  porcelain  crucible,  and  heat  until  the  white  fumes 
of  ammonium  salts  cease  to  be  given  off.  Dissolve  the 
residue  in  a  few  drops  of  H2O  with  a  drop  or  two  of  HC1, 
and  add  PtCl4 ;  K,  if  present,  will  be  precipitated. 

A  loop  of  platinum  wire  dipped  in  the  original  solution 
and  held  in  the  colorless  gas  flame  will  give  evidence  of 
Na  and  Li. 

If  Na  be  present  in  excess,  so  as  to  obscure  the  Li  flame, 
evaporate  a  portion  of  the  original  solution  to  dryness,  dis- 
solve in  the  smallest  possible  amount  of  H2O,  add  Na2HPO4 
and  boil,  filter  off  the  Li3PO4,  wash  with  a  little  hot  water 
containing  NH4OH,  and  dissolve  in  a  few  drops  of  HC1. 
With  this  solution  Li  may  be  detected  by  the  flame  test. 


GROUP    II.— BARIUM,    STRONTIUM,    CALCIUM, 
MAGNESIUM. 

REACTIONS   OF   BARIUM  (Ba). 
Use  a  solution  of  barium  chloride  (BaCl2). 
I.  H2SO4   produces   an    immediate   precipitate  of  barium 
sulphate — BaSO4,  insoluble  in  boiling  hydrochloric  or  nitric 

acid. 

BaCl2  +  H2SO4  =  BaSO4  -f  2HC1. 


36  ANALYTICAL   CHEMISTRY. 

2.  K2CrO4,  even  in  dilute  solutions,  causes  a  yellow  precipi- 
tate of  barium  chromate — BaCrO4,  soluble  in  hydrochloric 
or  nitric  acid,  but  insoluble  in  acetic  acid. 

BaCl2  +  K2CrO4  =  BaCrO4  +  2KC1. 

3.  (NH4)2CO3  precipitates  white  barium  carbonate — BaCO3, 
soluble  in  acetic  acid. 

BaCl2  +  (NH4)2CO3  =  BaCO3  -f  2NH4C1. 

4.  (NH4)2HPO4   produces   a  white   precipitate  of  barium 
phosphate — BaHPO4,  soluble  in  acetic  and  in  hydrochloric 

acid. 

BaCl2  +  (NH4)2HP04  =  BaHPO4  +  2NH4C1. 

5.  (NH4)2C2O4  causes  a  white  precipitate  of  barium  oxalate 
— BaC2O4,  slightly  soluble  in  acetic  acid.     This  precipitation 
will  not  take  place  in  very  dilute  solutions. 

BaCl2  +  (NHt)2C2O4  r=  BaC2O4  +  2NH4C1. 

6.  A  loop  of  platinum  wire  moistened  with  the  solution 
colors  the  gas  flame  green  when  held  in  it. 

REACTIONS   OF   STRONTIUM  (Sr). 
Use  a  solution  of  strontium  nitrate  (Sr(NO3)2). 

1.  H2SO4  forms  a  white  precipitate  of  strontium  sulphate 
— SrSO4,  immediately,  if  the  solution  be  strong,  but  not  until 
after  some  time,  if  it  be  very  dilute. 

2.  K2CrO4  produces  no  precipitate  in  the  presence  of  acetic 
acid,  but  if  the  solution  be  made  alkaline  with  KOH,  a  yellow 
precipitate,  strontium  chromate — SrCrO4,  falls. 

3.  (NH4)2CO3  produces  a  white  precipitate  of  strontium 
carbonate — SrCO3,  soluble  in  acetic  and  the  stronger  acids. 
Na2CO3  produces  the  same  precipitate. 

4.  (NH4)2HPO4   forms   a   white   precipitate    of  strontium 
phosphate — SrHPO4,  soluble  in  acids. 

5.  (NH4)2C2O4  causes  the  precipitation  of  white  strontium 
oxalate — SrC2O4,  sparingly  soluble  in  acetic  acid,  but  readily 
soluble  in  HC1. 

6.  Strontium  salts  impart  an  intense  red  to  the  colorless  gas 
flame. 


BASES.  37 

REACTIONS  OF  CALCIUM  (Ca). 
Use  a  solution  of  calcium  chloride  (CaCl2). 

1.  H2SO4,  in   moderately   dilute    solutions,  forms  a  white 
precipitate  of  calcium  sulphate — CaSO4,  soluble  in  excess  of 
water. 

2.  (NH4)2CO3  or   Na2CO3  produces  a  white  precipitate  of 
calcium  carbonate — CaCO3,  soluble  in  acids.     This  precipi- 
tation is  not  complete  unless  the  solution  is  boiled. 

3.  (NH4)2HPO4  causes  the  precipitation  of  calcium  phos- 
phate— CaHPO4,  soluble  in  acetic  and  the  stronger  acids. 

4.  (NH4)2C2O4  produces  a  white  precipitate  of  calcium  oxa- 
late — CaC2O4,  insoluble  in  acetic  acid,  soluble  in  hydrochloric 
or  nitric  acid. 

5.  The  salts  of  calcium  color  the  flame  yellowish-red. 

REACTIONS  OF  MAGNESIUM  (Mg). 
Use  a  solution  of  magnesium  sulphate  (MgSO4). 

1.  (NH4)2CO3  forms  a  white  precipitate  of  magnesium-am- 
monium carbonate— MgCO3(NH4)2CO3,  soluble  in  NH4C1. 
By  preceding  the  addition  of  the  reagent  by  that  of  NH4C1, 
a  much  smaller  quantity  will  suffice  to  keep  the  precipitate 
in    solution   than  will  be   required  to   dissolve   it  after  once 
formed. 

MgS04  +  2(NH4)2C03  =  MgC03(NH4)2C03  +  (NH4)2SO4. 

2.  KOH,NaOH  or  NH4OH  produces  a  white  precipitate  of 
magnesium  hydrate — Mg(OH)2,  soluble  in  NH4C1. 

3.  (NH4)2HPO4  with  NH4C1  and  NH4OH  produces  a  white 
crystalline  precipitate  of  ammonium-magnesium  phosphate 
— Mg(NH4)PO4,  slightly  soluble  in  water,  but  almost  insoluble 
in  water  containing  NH4OH.     Violent   agitation   or  stirring 
assists   in  the   formation   of  this   precipitate.     (NH4)2HAsO4 
under  similar  circumstances  precipitates  white  Mg(NH4)AsO4. 

MgSO4  +  NH4OH  +  Na2HPO4  =  Mg(NH)4PO4  +  Na2SO4  -f  H2O. 

The  ammonium  chloride  takes  no  part  in  the  reaction,  except 
to  keep  magnesium  hydrate  from  precipitating. 

4.  Magnesium  salts  impart  no  color  to  the  flame. 


38  ANALYTICAL    CHEMISTRY. 

SUMMARY  OF  TESTS  WITH  SOLUBLE  SALTS  OF  GROUP  II. 


Ba 

Sr 

Ca 

Mg 

H2S04 

White  Precipitate 
insoluble  in  acids 

White  Precipitate 
insoluble  in  acids 

White  Precipitate 
soluble  in  excess 
of  H2O 

No  Precipitate 

K2CrO4 

Yellow  Precipitate 
insoluble  in  acetic 
acid 

No  Precipitate 
unless  alkaline 

No  Precipitate 

No  Precipitate 

(NH4)2C03 

White  Precipitate 

White  Precipitate 

White  Precipitate 

White  Precipitate 
soluble  in  NH4C1 

(NH4)2HP04 

White  Precipitate 

White  Precipitate 

White  Precipitate 

White  Precipitate 

(NH4)2C204 

White  Precipitate 
in  strong  solution 

White  Precipitate 
in  strong  solution 

White  Precipitate 
in  dilute  solution 

No  Precipitate 
unless  concentrated 

NH4OH- 

No  Precipitate 

No  Precipitate 

No  Precipitate 

White  Precipitate 

DIRECTIONS  FOR  THE  DETECTION  OF  THE  BASES   IN  A  SOLUTION 

CONTAINING  SOLUBLE  SALTS  OF  GROUP  II. 

Add  NH4C1,  NH4OH  and  (NH4)2CO3,  boil  and  filter. 


Ppt   Ba,  Sr,  Ca. 
Wash,  dissolve  in  HC2H3O2,  add  K2CrO4, 
filter. 

Filt.  Mg. 
Add  (NH4)2HPO4,  agitate.     White  ppt. 
if  Mg  be  present. 

Ppt.  Ba. 
Yellow. 

Filtrate  Sr,  Ca. 
Add  very  dilute  H2SO4,  allow  to 
stand  10  minutes,  filter. 

Ppt.  Sr. 
Confirm  by 
flame  test. 

Filt.  Ca. 
Add  NH4OH  and 
(NH4)2C204, 
white  ppt. 

In  order  to  thoroughly  acquaint  the  student  with  the  method 
of  analysis  by  this  and  subsequent  charts,  the  following  expla- 
nation is  given,  in  the  belief  that  a  careful  study  of  it,  until 
perfectly  understood,  will  enable  the  student  to  follow  all  the 
charts  which  come  after,  and  which  are  simply  enlargements 
of  this  scheme. 

To  a  small  quantity  of  the  solution  in  a  test  tube  add  an 
equal  volume  of  solution  of  NH4C1,  close  and  invert  the  tube, 
so  as  to  thoroughly  mix  the  contents,  add  NH4OH  until,  after 
mixing,  the  solution  smells  distinctly  of  it,  and  then  add 
(NH4)2CO3,  boil  and  filter. 

We  suppose  all  four  of  the  bases  under  consideration  to  be 
present  until  their  absence  is  proven ;  so  by  this  treatment  we 
divide  them  into  two  groups.  The  precipitate  consists  of  Ba, 
Sr  and  Ca,  and  the  filtrate  of  Mg.  A  little  more  (NH4)2  CO3 


BASES. 


39 


should  be  added  to  this  filtrate  and  the  whole  again  boiled, 
to  make  sure  that  all  the  insoluble  carbonates  have  been 
precipitated.  If  this  is  found  to  be  the  case,  (NH4\HPO4 
is  added  to  the  filtrate,  and  Mg,  if  present,  will  form  a  white 
precipitate. 

The  first  precipitate,  consisting  of  Ba,  Sr  and  Ca,  having,  in 
the  meantime,  been  washed  by  forcing  a  jet  of  water  from  the 
wash  bottle  on  it,  is  dissolved  in  HC2H3O2,  and  to  the  solu- 
tion K2CrO4  added  ;  this  again  divides  the  solution  into  a  pre- 
cipitate and  a  filtrate.  The  yellow  precipitate  represents  Ba, 
while  the  filtrate  contains  the  Sr  and  Ca.  To  this  filtrate  very 
dilute  H2SO4  (made  by  adding  a  small  quantity  of  H2SO4  to 
about  ten  times  its  volume  of  water)  is  added,  and  the  mix- 
ture allowed  to  stand  ten  minutes,  for  the  SrSO4  to  form  and 
subside,  then  filtered,  which  again  gives  us  a  precipitate,  repre- 
senting Sr,  and  a  filtrate  indicating,  after  the  addition  of 
NH4OH  and  (NH4\C2O4,  the  presence  or  absence  of  Ca. 

DIRECTIONS   FOR  THE  ANALYSIS  OF  A  SOLUTION  CONTAINING 

SOLUBLE  SALTS  OF  ALL  THE  PRECEDING  ELEMENTS. 

Add  NH4C1,  NH4OH  and  (NH4)2CO3,  boil  and  filter. 


Ppt.  Ba,  Sr,  Ca. 
Wash,  dissolve  in  HC,H,O2,  add  K2CrO4, 
filter. 

Filtrate  Mg,  K,  Na,  Li,  NH4. 
Add  (NH4)2HPO4,  agitate,  filter. 

SJ: 

Filtrate  K,  Na,  Li,  NH4. 
Evaporate  to  dryness,  ignite,  dissolve 
in  a  small  quantity  of  H,O,  add 
Na2HP04,  boil,  filter. 

Ppt.  Ba. 
Yellow. 

Filtrate  Sr,  Ca. 
Add  very  dilute  H2SO4,  allow  to 
stand,  filter. 

Ppt.  Sr. 
Confirm  by 
flame  test. 

Filt.  Ca. 
Add  NH4OH  and 
(NH4)2C204, 
white  ppt. 

* 

Confirm 
by  flame 
test. 

Filtrate  K,  Na,  NH4. 
Concentrate,  add  HC1  and 
PtCl4  yellow  ppt.  —  K. 
Test  for  Na  and  NH4 
in  original  solution. 

PRECAUTIONS  TO  BE  OBSERVED  IN  THE  PRECEDING  CHARTS. 
Ammonium  chloride  must  be  added  in  excess,  in  order  to 
keep  the  magnesium  salts  in  solution  when  the  hydrate  and 
carbonate  are  added.  Ammonium  hydrate  is  added  until 
the  liquid  smells  of  it.  Ammonium  carbonate  is  added  as 
long  as  a  precipitate  is  produced.  In  order  to  determine  this 
to  a  certainty,  a  portion  of  the  filtrate  is  tested  with  a  little 
more  of  the  reagent,  when,  if  no  precipitate  occurs,  the  analysis 
may  be  proceeded  with.  This  precaution  of  applying  more  of 


40  ANALYTICAL   CHEMISTRY. 

the  reagent  to  a  portion  of  the  filtrate,  to  prove  the  complete 
precipitation,  should  be  exercised  in  every  case,  as  it  is  im- 
portant to  add  just  sufficient  of  the  reagent  to  accomplish  the 
object,  but  always  to  avoid  a  large  excess. 

The  analysis  of  the  above  solutions  may  be  much  simplified 
in  many  cases  by  adding  a  solution  of  CaSO4  to  the  original 
solution.  If  a  precipitate  form  immediately,  Ba  is  present, 
Sr  and  Ca  may  be.  If  a  precipitate  form  after  some  time, 
Ba  is  absent,  Sr  is  present,  and  Ca  may  be  present.  If  no  pre- 
cipitate be  formed,  Ba  and  Sr  are  absent,  and  Ca  may  be  tested 
for  in  another  portion  with  (NH4)2C2O4.  The  above  charts 
suppose  all  the  elements  to  be  present,  in  which  case  they 
afford  the  simplest  means  of  detecting  them. 


GROUP  III.— MANGANESE,  ZINC,  COBALT, 
NICKEL. 

REACTIONS    OF    MANGANESE   (Mn). 
Use  a  solution  ofmanganous  sulphate  (MnSO4). 

1.  NH4HS,  in  neutral  or  alkaline  solution,  precipitates  the 
flesh-colored  manganous  sulphide — MnS,  which,  on  exposure 
to  air,  becomes  brown.     HC1,  HNO3  and  HC2H3O2  dissolve 
this  precipitate,  but  it  is  insoluble  in  alkalies.     NH4C1  facili- 
tates the  separation  of  the  precipitate,  while  the  salts  of  the 
organic  acids  and  excess  of  NH4OH  prevent  it. 

2.  KOH  or  NaOH  produces  a  whitish  precipitate  of  man- 
ganous hydrate — Mn(OH)2,  insoluble  in  excess. 

MnSO4  -j-  2KOH  =Mn(OH;2  +  K2SO4. 

3.  NH4OH   likewise  precipitates  Mn(OH)2,  first  white,  but 
becoming  rapidly  brown,  partly  soluble  in  excess.     This  pre- 
cipitation is  prevented  by  the  previous  addition  of  NH4C1. 

4.  (NH4)2CO3  produces  a  white  precipitate  of  manganous 
carbonate — MnCO3,  insoluble  in  excess. 

5.  HNO3  and  red  oxide  of  lead  will,  on  heating  and  allow- 
ing the  precipitate  to  subside,  impart  to  the  supernatant  liquid 
a  red  color,  due  to  permanganic  acid — H2Mn2O8.     Hydro- 
chloric acid  and  chlorides  interfere  with  this  reaction.     This 
is  known  as  Crum's  process  for  detecting  manganese. 


BASES.  41 

6.  A  fragment  of  a  manganese  salt  fused  on  platinum  foil 
with  K2CO3  and  KNO3  will  form  a  green  mass  containing 
potassium  manganate — K2MnO4. 

/.  A  borax  bead  (formed  by  fusing  on  the  loop  of  a  plati- 
num wire  some  borax  until  it  becomes  a  clear  glass)  with 
manganese,  in  the  oxidizing  blowpipe  flame,  becomes  violet 
while  hot,  and  a  fine  amethyst  color  on  cooling. 

REACTIONS   OF   ZINC  (Zn). 
Use  a  solution  of  zinc  sulphate  (ZnSO4). 

1.  NH4HS  produces  a  white  precipitate  of  zinc  sulphide — 
ZnS,  insoluble  in  acetic  acid,  readily  soluble  in  dilute  hydro- 
chloric acid. 

2.  KOH,  NaOH   and   NH4OH   give  white  precipitates  of 
zinc  hydrate — Zn(OH)2,  readily  soluble   in  excess,  forming 
zincates  as  Zn(OK)2.     Zn(OH)2  is  again  precipitated  on  boil- 
ing. 

3.  (NH4)2CO3  forms  a  white  precipitate  of  basic  zinc  car- 
bonate— (ZnCO3)2(Zn(OH)2)3,  readily  soluble  in  excess. 

5(NH4)2C03  -f  5ZnS04  +  3H2O  = 
2ZnC033Zn(OH)2  +  5(NH4)2SO4  +  3CO2. 

4.  K2CO3  or  Na2CO3  produces  a  similar  precipitate,  insoluble 
in  excess. 

5.  On  charcoal,  before  the  blowpipe,  metallic  zinc  volatilizes 
and  burns,  forming  an  incrustation  of  oxide,  which  is  yellow 
while  hot,  becoming  white   on    cooling ;    if  this  coating   be 
moistened  with  a  drop  of  cobaltous  nitrate,  and  again  heated 
in  the  outer  flame,  it  becomes  green. 

REACTIONS   OF   COBALT  (Co). 
Use  a  solution  of  cobaltous  nitrate  (Co  (NO3)2). 

i.  NH4HS  produces  a  black  precipitate  of  cobaltous  sul- 
phide— CoS,  insoluble  in  acetic  acid  and  cold  dilute  hydro- 
chloric acid.  The  precipitation  is  promoted  by  the  presence 
of  NH4C1. 

2NH4HS  +  Co(NO3)2  =  CoS  +  2NH4NO8  -f  H2S. 


42  ANALYTICAL    CHEMISTRY. 

2.  KOH  or  NaOH  produces  a  blue  precipitate  of  cobaltous 
hydrate — Co(OH)2,  insoluble  in  excess,  and  becoming  pink 
on  boiling  or  exposure  to  air. 

3.  NH4OH  causes  a  similar  precipitate  of  Co(OH)2,  soluble 
in  excess  with  a  red  color.     Sugar  and  some  other  organic 
compounds  prevent  the  precipitations  by  the  alkalies.     The 
alkaline  carbonates  behave  like  their  respective  hydrates. 

4.  KCN  gives  a  red-brown  precipitate  of  cobaltous  cyanide 
— Co(CN)2,  soluble  in  excess  and  reprecipitated  by  HC1 ;  if, 
however,  the  solution  be  boiled  with  only  a  few  drops  of  HC1, 
the  cobaltous  cyanide  will  not  be  precipitated  on  the  further 
addition  of  HC1,  on  account  of  the  formation   of  potassium 
cobalti-cyanide — K6Co2(CN)12.    This  experiment  should  be  per- 
formed in  a  fume  closet,  in  order  to  avoid  inhaling  the  fumes  of 
hydrocyanic  acid. 

5.  Salts  of  cobalt  color   the   borax  bead  blue  before  the 
blowpipe. 

REACTIONS   OF   NICKEL  (Ni). 
Use  a  solution  of  nickelous  sulphate  (NiSO4). 

1.  NH4HS  forms  a  black  precipitate  of  nickelous  sulphide 
— NiS,  insoluble  in   acetic  acid  and  cold  dilute  hydrochloric 
acid.     The    precipitation    is    promoted    by   the   presence    of 
NH4C1. 

2.  KOH  or  NaOH  produces  a  green  precipitate  of  nickel- 
ous hydrate — Ni(OH)2,  insoluble  in  excess. 

3.  NH4OH  gives  a  similar  precipitate,  soluble  in  excess,  with 
a  blue  color.     Sugar  and  some  other  organic  compounds  pre- 
vent the  precipitation  by  the  alkalies.     The  alkaline  carbon- 
ates behave  like  their  respective  hydrates. 

4.  KCN  produces  a  yellowish-green  precipitate  of  nickel- 
ous cyanide — Ni(CN)2,  soluble  in  excess,  and  re-precipitated 
by  HC1  even  after  boiling,  also  precipitated,  after  adding  HCl 
and  boiling,  by  KOH.     This  is   used  as   a   method  of  dis- 
tinguishing and   separating  nickel  and  cobalt,  the  latter  not 
precipitating   under   these   circumstances   with    KOH.     This 
experiment  should  be  performed  in  a  fume  closet. 


BASES. 


43 


5.  The  salts  of  nickel  color  the  borax  bead  violet  while  hot, 
and  reddish-brown  when  cold. 


SUMMARY  OF  TESTS  WITH  SOLUBLE  SALTS  OF  GROUP   III. 


Mn 

Zn 

Co 

Ni 

NH4HS 

Flesh  colored 
precipitate. 

White  precipitate 

Black  precipitate 

Black  precipitate 

KOH 

White  precipitate 
insoluble  in  excess 

White  precipitate 
soluble  in  excess 

Blue  precipitate 
insoluble  in  excess 

Green  precipitate 
insoluble  in  excess 

NH4OH 

White  precipitate 
soluble  in  excess 

White  precipitate 
soluble  in  excess 

Blue  precipitate 
soluble  in  excess 

Green  precipitate 
soluble  in  excess 

(NH4)2C03 

White  precipitate 
insoluble  in  excess 

White  precipitate 
soluble  in  excess 

Blue  precipitate 
soluble  in  excess 

Green  precipitate 
soluble  in  excess 

DIRECTIONS   FOR  THE  DETECTION  OF  THE  BASES  IN  A  SOLUTION 
CONTAINING  SOLUBLE  SALTS  OF  GROUP   III. 

Mn,  Zn,  Co,  Ni.     Add  NH4OH,  then  acidify  slightly  with  HC2H3O2  and  add  H2S  until  the 
liquid  smells  strongly ;  falter. 


Precipitate  Zn,  Co,  Ni. 

Wash,  dissolve  in  HC1  and  HNO3,  evaporate  excess  of  acid, 
add  KOH  in  excess,  filter. 


Precip.  Co,  Ni. 

Wash,  dissolve  in  HC1,  add  KCN  in  excess  and 

boil  with  HC1  until  all  odor  of  HCN  disappears, 

then  add  KOH  in  excess,  filter. 


Precip.  Ni. 
Green. 


Filt.  Co. 

Evaporate  to  dryness  and  test 
with  borax  bead. 


Filt.  Zn. 

Add  NH4HS 

white  i  pt. 


Filtrate  Mn. 

Add  NH4OH  and  NH4HS, 
ppt.  pink. 


44 


ANALYTICAL    CHEMISTRY. 


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FOR  THE  DETECT 

15 

Ppt.  Zn,  Co,  Ni. 
n  HCI  and  HNO3,  evapc 
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BASES.  45 

GROUP  IV.— IRON,  CERIUM,  ALUMINIUM, 
CHROMIUM. 

REACTIONS  OF  IRON  in  ferrous  state  (Fe)11. 
Use  a  solution  of  ferrous  sulphate  (FeSO4). 

1.  K4Fe(CN)6/in  a  neutral  or  acid  solution,  gives  a  white 
(rapidly  changing   to    light   blue)   precipitate    of  potassium 
ferrous  ferrocyanide — K2Fe2(CN)6,  also  known  as  Everett's 
salt.     Alkalies  decompose  this  precipitate,  forming  ferrous  hy- 
drate and  a  ferrocyanide  of  the  base  used. 

K4Fe(CN)6  +  FeS04  =  K2Fe2(CN)6  +  K2SO4. 

2.  K6Fe2(CN)12,  in  a  neutral  or  slightly  acid  solution,  forms 
a  dark  blue  precipitate  of  ferrous  ferricyanide — Fe3Fe2(CN)12, 
known  as  Turnbull's  blue.     If  the  solutions  be  very  dilute 
there  is  produced  merely  a  deep  blue-green  coloration. 

K6Fe2(CN)12  +  3FeS04  =  Fe3Fe2(CN)12  +  3K2SO4. 

3.  KCNS  produces  no  change. 

4.  H2S  in  acid  solution  does  not  form  a  precipitate. 

5.  NH4HS,  with  a  neutral  or  alkaline  solution,  forms  a  black 
precipitate  of  ferrous  sulphide,  soluble  in  HC1  or  HNO3. 
NH4C1  promotes  the  formation  of  this  precipitate. 

6.  NH4OH  in  the  absence  of  NH4C1  produces  a  dirty  green 
precipitate   of  ferrous  hydrate — Fe(OH)2.     This  precipitate 
rapidly  becomes  reddish-brown,  owing  to  absorption  of  oxygen. 

7.  KOH  produces  a  dirty  green  precipitate  of  ferrous  hy- 
drate— Fe(OH)2,  similar  to  that  produced  by  ammonia.     Non- 
volatile organic  substances,  as  sugar  and  some  acids,  retard 
the  precipitations  by  NH4OH  and  KOH. 

8.  Na2CO3  causes  a  white  precipitate  of  ferrous  carbonate 
— FeCO3>  which  rapidly  becomes  brown,  from  absorption  of 
oxygen.     This  rapid  oxidation  is  prevented  by  the  use  of  dis- 
tilled water  and  sugar. 

REACTIONS  OF  IRON  in  ferric  state  (Fe2)VI. 

Use  a  solution  of  ferric  chloride  (Fe2Cl6). 
i.  K4Fe(CN)6,  in  a  neutral  or  acid  solution,  produces  a  dark 
blue  precipitate  of  ferric  ferrocyanide — (Fe2)2(Fe(CN)6)3,  de- 
composed by  alkalies. 

3K4Fe(CN)6  +  2Fe2Cl6  i=  (Fe2)2(Fe(CN)6)3  +  12  KC1. 


46  ANALYTICAL    CHEMISTRY. 

2.  K6Fe2(CN)12  forms  no  precipitate,  but  produces  a  deep 
reddish-brown  color.     The  olive-green  color  sometimes  pro- 
duced in  this  reaction  is  due  to  traces  of  ferrous  salt. 

3.  KCNS  imparts  to  acid  solutions  a  deep  blood-red  color, 
due  to  the  formation  of  ferric  sulphocyanate.     This  color  is 
immediately  destroyed  by  HgG2.     Dilute  solutions  show  this 
reaction  best. 

4.  H2S   forms   in   acid  solutions   a  white  turbidity  due  to 
separation  of  sulphur ;  the  ferric  salt  being  at  the  same  time 
reduced  to  the  ferrous  condition. 

2Fe2Cl6  +  2H2S  ==  4FeCl2  +  4HC1  +  S2. 

In    alkaline    solutions    this    reagent    acts    as    an    alkaline 
sulphide. 

5.  NH4HS  causes  a  black  precipitate  of  ferrous  sulphide 
— FeS,  sulphur  separating  at  the  same  time. 

6.  NH4OH    precipitates    reddish-brown    ferric  hydrate  — 
Fe2(OH)6;  non-volatile  organic  acids  and  sugar  prevent  this 
precipitation. 

7.  KOH  and  NaOH  react  like  NH4OH. 

8.  Na2CO3  and   the   other   alkaline   carbonates   precipitate 
ferric  hydrate— Fe,(OH)6. 

9.  With  borax  in  the  oxidizing  blowpipe  flame,  ferrous  and 
ferric  compounds  give  dark  yellow  to  red  colored  beads  while 
hot,  and  yellow  when  cold.     In  the  reducing  flame  the  beads 
change  to  bottle-green. 

REACTIONS  OF  CERIUM  (Ce). 
Use  a  solution  of  cerous  chloride  (Ce2Cl6). 

1.  NH4HS  causes  a  white  precipitate  of  cerous  hydrate — 
Ce2(OH)6. 

2.  NH4OH  produces  the  same  white  precipitate  of  Ce2(OH)6, 
as  also  do  KOH  and  NaOH. 

3.  (NH4)2C2O4  or  H2C2O4  forms  a  white  precipitate  of  cerous 
oxalate — Ce2(C2O4)3.     Organic  matter  does  not  interfere  with 
the  formation  of  this  precipitate. 

4.  With  a  borax  bead  before  the  blowpipe  the  salts  of  cerium 
behave  like  those  of  iron. 


BASES.  47 

REACTIONS   OF 'ALUMINIUM  (Al). 
Use  a  solution  of  alum  (K2SO4A12(SO4)3). 

1.  NH4HS  produces  a  white  precipitate  of  aluminium  hy- 
drate— A12(OH)6,  while  H2S  escapes. 

6NH4HS  +  K2S04A12(S04)3  +  6H2O  =  A12(OH)6  +  3(NH4)2SO4 
-f  K2SO4  +  6H2S. 

2.  NH4OH  forms  a  white  precipitate  of  aluminium  hydrate 
— A12(OH)6,  insoluble  in  excess ;  this  is  an  important  distinc- 
tion from  zinc. 

3.  KOH  and  NaOH  produce  a  similar  white  precipitate  of 
A12(OH)6,  soluble  in  excess,  forming  aluminates  of  the  base, 
as  A12(OK)6.     This  is  not  reprecipitated  by  boiling  (distinction 
from  chromium). 

4.  Na2CO3   and   the    other   alkaline    carbonates  precipitate 
white  gelatinous  aluminium  hydrate — A12(OH)6,  insoluble  in 
excess,  CO2  escaping  at  the  same  time. 

3Na2C03  +  K2S04A12(S04)3  -f  3H2O  =  A12(OH)6  +  3Na2SO4 
+  K2S04  -j-  3C02. 

The  presence  of  non-volatile  organic  acids  and  sugar  prevent 
the  complete  precipitation  in  the  above  reactions. 

REACTIONS  OF  CHROMIUM  (Cr). 
Use  a  solution  of  chromic  chloride  (Cr2Cl6). 

1.  NH4HS  produces  a  greenish  precipitate  of  chromic  hy- 
drate— Cr2(OH)6,  H2S  escaping. 

2.  NH4OH  forms  the  same  greenish  precipitate  of  Cr2(OH)6, 
insoluble  in  excess. 

3.  KOH  and  NaOH  produce  the  same  precipitate  soluble 
in  excess,  but  reprecipitated  on  boiling  (distinction  from  alu- 
minium). 

4.  Na2CO3  and   the   other   alkaline   carbonates,  precipitate 
green  basic  carbonates. 

5.  KNO3  and  K2CO3  fused  with  chromium  compounds  be- 
come yellow  from  formation  of  potassium  chromate — K2CrO4. 
AgNO3  and  Pb(C2H3O2)2  are  important  tests  for  this  compound ; 
the  former  precipitates  red  silver  chromate,  the  latter  yellow 
lead  chromate. 

6.  With  the  borax  bead  in  the  inner  blowpipe  flame,  chro- 
mium compounds  give  a  green  color. 


48  ANALYTICAL   CHEMISTRY. 

SUMMARY  OF  TESTS  WITH  SOLUBLE  SALTS  OF  GROUP   IV. 


Fe(ous) 

Fe(ic) 

Ce 

Al 

Cr 

K4Fe(CN)6 

White  ppt. 
turning  blue. 

Deep  blue 
precipitate. 

K6Fe2(CN)12 

Deep  blue 
precipitate. 

No  ppt. 
brownish-red 
color. 

KCNS 

No  change. 

Blood-red 
color. 

NH4HS 

Black  ppt. 

Bk  ck  ppt. 

White  ppt. 

White  ppt. 

Greenish  ppt. 

NH4OH 

Dirty  green 
ppt. 

Reddish- 
brown  ppt. 

White  ppt. 

White  ppt.  in- 
soluble in  excess. 

Greenish  ppt.  in- 
soluble in  excess. 

KOH 

Dirty  green 
ppt. 

Reddish- 
brown  ppt. 

White  ppt. 

White  ppt.  solu- 
ble in  excess,  not 
reprecipitated 
by  boiling. 

Green  ppt. 

soluble  in  excess, 
reprecipitated 
by  boiling. 

Na2C03" 

White  ppt.  be- 
coming dark. 

Reddish- 
brown  ppt. 

White  ppt. 

White  ppt. 

Green  ppt. 

DIRECTIONS  FOR  THE  DETECTION  OF  THE  BASES  IN  A  SOLUTION 
CONTAINING  SOLUBLE  SALTS  OF  GROUP  IV. 

Evaporate  a  portion  of  the  solution  to  dryness,  fuse  on  platinum  foil  with  Na2CO3  and 
KNO3,  boil  with  water  and  filter. 


Residue,  Fe,  Ce. 
Dissolve  in  concentrated  H2SO4  and  C2H5OH, 
divide  in  two  parts. 

Filtrate,  Al,  Cr. 
Yellow  if  Cr  be  present,  divide  in  two  parts. 

Fe. 
Dilute,  test  with 
K4Fe(CN)6. 
Test  original  solution 
with  K4Fe(CN)6 
and  K6Fe2(CN)12  for 
Fe(ous)  and  Fe(ic). 

Ce. 
Dilute,  add 
H3C6H607,  NH4OH 
and  Na2HP04, 
white  ppt. 

Al. 
Add  NH4C1  and  warm, 
white  gelatinous 
ppt. 

Cr. 
Add  HC,H,O2  and 
Pb(C2H302)2 
yellow  ppt. 

BASES. 


49 


4 


ll 


SB.:? 


U  ^ 


s 


^ 

6°, 


^ 
^| 


•a  « 


offi^i 


Filt.  M 
ff  excess  o 
HC2H302 


Filt. 
K,  Na,  Li,  NH4. 

Evaporate  to  dryness 
ignite,  dissolve  in 


ff     o 
STJJ5 


rt<  — 


E    fc 


1 

«  c 


SB 


-a    3 


.        O 

&* 


50  ANALYTICAL   CHEMISTRY. 

PRECAUTIONS  AND  OBSERVATIONS  ON  THE  PRECEDING  CHART. 

1.  It  is  particularly  desirable  not  to  filter  immediately  after 
adding  NH4HS  in  the  first  precipitation ;  otherwise,  the  Mn 
will  not  be  thoroughly  precipitated. 

2.  When  the  first  precipitate  is  dissolved  in  HC1  and  HNO3, 
great  care  must  be  exercised  to  thoroughly  oxidize  the  Fe  by 
boiling  with  the  HNO3,  otherwise,  the  precipitation  by  NH4OH 
will  not  be  complete. 

3.  The   precipitate    of   aluminium    hydrate    is    difficult    to 
observe,    as    it   floats    in   the    solution    instead    of  falling   to 
the  bottom.     Warming  the   solution  will    usually  render    it 
visible. 

4.  When  phosphoric  acid  is  present,  the  members  of  Group 
II  precipitate  with  Group  IV,  and  require  an  entirely  different 
method  of  separation.     It  is  desirable,  however,  for  the  student 
first  to  familiarize  himself  with  these  simpler  soluble  salts,  and 
undertake  the  more  difficult  cases  of  salts  insoluble  in  water 
after  the  acids  have  been  considered.     See  page  83. 


GROUP  V.— ARSENIC,  ANTIMONY,  TIN,  GOLD, 
PLATINUM. 

REACTIONS   OF  ARSENIC   (As). 

(a)    ARSENIOUS   COMPOUNDS. 

Use  a  solution  of  As2O3,  in  water. 

1.  H2S  passed  into  the  solution  produces  a  yellow  color, 
but  no  precipitate  until  HC1  is  added,  when  a  yellow  precipi- 
tate of  arsenious  sulphide — As2S3,  falls.     This  precipitate  is 
insoluble  in  strong  HC1,  but  soluble  in  NH4HS,  NH4OH  and 
(NH4)2C03. 

As203  +  3H2S  =  As2S3  -f  3H20. 

2.  NH4HS  causes  the  formation  of  arsenious  sulphide, 
which  remains   in  solution  as  ammonium   sulpharsenite — 
(NH4)3AsS3.     On  the  addition  of  HC1  arsenious  sulphide  is 
precipitated. 

3.  AgNO3  produces    no   precipitate    until  a  few   drops   of 
dilute  ammonia  solution  are  added,  when  a  yellow  precipitate 
of  silver  arsenite — Ag3AsO3,  falls,  soluble  in  HNO3  and  in 
NH4OH. 


BASES.  51 

4.  CuSO4,  under  similar  circumstances,  produces  a  yellowish- 
green  precipitate  of  cupric  arsenite — CuHAsO3. 

(&)    ARSENIC    COMPOUNDS. 

Use  a  solution  of  sodium  arsenate  (Na2HAsO4). 

1.  H2S  causes,  in  acid  solution  only,  a  yellow  precipitate  of 
arseniou  ssulphide — As2S3,  mixed  with    sulphur.     This  re- 
action takes  place  slowly,  but  is  accelerated  by  heat. 

5H2S  4-  2Na2HAsO4  +  4HC1  =  As2S3  +  S2  -f-  4NaCl  4-  8H2O. 

2.  NH4HS   produces    no   precipitate,   but    forms    arsenic 
sulphide — As2S5,  which  remains  in  solution  as  ammonium 
sulpharsenate — (NH4)3AsS4.      Upon    the    addition    of    HC1 
As2S5  is  precipitated,  and  not  As2S3  and  S. 

3.  AgNO3,  with  a  small  amount  NH4OH,  produces  a  choco- 
late-colored precipitate  of  silver  arsenate — Ag3AsO4,  soluble 
in  HNO3  and  NH4OH. 

4.  CuSO4,  under  similar  circumstances,  forms  a  bluish-green 
precipitate  of  cupric  arsenate — CuHAsO4. 

The  following  tests   are  applicable  to  both  arsenious  and 
arsenic  compounds. 

1.  Mars/is   Test. — Generate    hydrogen    in   the    usual  way, 
allowing  it  to  escape  through  a  glass  tube  drawn  out  at  the 
end  so  as  form  a  small  orifice  (Fig.  4).     In  very  exact  cases, 
the  gas   should  be  dried   by  passing  over  calcium   chloride. 
When  all  the  air  has  been  expelled  (which  should  be  deter- 
mined   by  collecting  a  small    test  tube  full  and  holding  its 
mouth  to  a  flame;  if  the  gas  burn  quietly,  without  explosion, 
it  is  pure),  ignite  the  escaping  gas  ;    it  should  burn  with  a 
colorless   or  yellow  flame ;  in  the  latter  case  it  •  is  due  to  the 
sodium  in  the  glass.     A  piece  of  cold  porcelain — a  small  cru- 
cible lid  is  best — is  pressed  down  on  the  flame ;  there  should 
be  no  deposit  on  it.     Add  now,  through  the  funnel  tube,  a 
solution  of  arsenic,  washing  it  down  with  a  little  water.     The 
flame  will  become  of  a  pale  blue  color,  due  to  the  formation 
of  hydrogen  arsenide — H3As.     On  bringing  the  crucible  lid 
into  the  flame  now,  a   blackish-brown   deposit,  with  metallic 
lustre,  will  form  on  it.     This  deposit  is  readily  soluble  in  a 
solution  of  sodium  or  calcium  hypochlorite. 

2.  ReinscJCs   Test. — Boil  some  strips  of  copper  with  dilute 


52 


ANALYTICAL   CHEMISTRY. 


HC1 ;  if  no  discoloration  of  the  copper  takes  place,  the 
arsenic  solution  may  be  added.  The  copper  immediately 
becomes  coated  with  an  iron-gray  metallic  film.  Pour  off 
the  liquid,  dry  the  copper  by  holding  it  in  the  lamp  with 
the  fingers  so  it  may  not  become  too  hot,  place  in  a  clean, 
dry,  narrow  test  tube,  and  heat  gently,  when  a  white  ring 
of  As2O3  will  form  on  the  tube  above  the  copper,  readily 
distinguished  by  the  characteristic  octahedral  shape  of  the 
crystals. 

Fleitmarfs  Test. — Generate  hydrogen  in  a  test  tube  with 


FIG.  4. 


zinc  and  solution  of  potassium  hydrate ;  moisten  a  piece  of 
filter  paper  with  one  drop  of  solution  of  silver  nitrate,  place 
it  over  the  mouth  of  the  tube  and  heat ;  there  should  be  no 
coloration  of  the  spot  on  the  paper.  Now  add  some  com- 
pound of  arsenic ;  the  silver  nitrate  will  immediately  become 
black,  owing  to  production  of  metallic  silver. 

H3As  +  (AgN03)6  +  (H20)3  =  H3As03  +  (HNO3)6  +  (Ag2)3. 

Before  the  blowpipe,  on  charcoal,  arsenic  volatilizes,  with 
the  characteristic  odor  of  garlic. 


BASES.  53 

REACTIONS   OF  ANTIMONY   (Sb). 
Use  a  solution  of  tartar  emetic  (KSbOC4H4O6),  acidified  with  HC1/ 

1.  H2S  forms  an  orange  precipitate  of  antimonous  sul- 
phide— Sb2S3,  soluble  in  NH4HS,  and  in  concentrated  HC1, 
but  insoluble  in  (NH4)2CO3. 

2.  NH4HS  produces  an  orange  precipitate  of  antimonous 
sulphide,    readily   soluble   in    excess,    forming   ammonium 
sulph-antimonite — (NH4)3SbS3,  from  which  HC1  again  pre- 
cipitates Sb2S3. 

3.  KOH   or  NaOH  precipitates  white,  bulky  antimonous 
hydrate — Sb(OH)3,  soluble  in  excess. 

4.  NH4OH  precipitates  the  same  compound  insoluble  in 
excess. 

5.  Marsh's  Test  gives  the  same  result  as  with  arsenic ;  the 
black  spot,  however,  is  insoluble  in  sodium  or  calcium  hypo- 
chlorite  solution,  but  soluble  in  NH4HS. 

6.  Reinsch's  Test  causes  a  deposit  on  copper,  as  with  arsenic, 
but  when  heated  in  a  tube  there  is  formed  a  white  amorphous 
ring,  which   is  readily  distinguished  from  the  crystalline  one 
of  arsenic. 

7.  Fleitman's    Test  gives   no   result   with    antimony   com- 
pounds. 

8.  On  charcoal,  with  Na2CO3,  before  the  blowpipe,  a  metallic 
globule  of  antimony  is  produced,  while   characteristic  fumes 
of  the  oxide  are  given  off. 

REACTIONS   OF   TIN   (Sn)v 

(a)    STANNOUS    COMPOUNDS. 

Use  a  solution  of  stannous  chloride  (SnCl2). 

1.  H2S  precipitates  dark  brown  stannous  sulphide — SnS, 
soluble   in  concentrated   HC1  and  in   (NH4)2S,  insoluble    in 
NH4HS. 

2.  KOH  or  NaOH  precipitates  white  stannous  hydrate — 
Sn(OH)2,  soluble  in  excess ;  on  boiling  this  solution  SnO  pre- 
cipitates. 

3.  NH4OH  precipitates  the  same   compound,  insoluble  in 
excess. 


54  ANALYTICAL   CHEMISTRY. 

4.  HgCl2  causes  a  white  precipitate  of  Hg2Q2,  converting 
the  SnCl2  into  SnCl4. 

2HgCl2  +  SnCl2  =  Hg2Cl2  +  SnQ4. 

This  precipitate  blackens  on  the  addition  of  NH4OH. 

(b)   STANNIC    COMPOUNDS. 

Use" a  solution  of  stannic  chloride  (SnCl4). 

1.  H2S  produces   a  yellow  precipitate  of  stannic  sulphide 
— SnS2,  soluble  in  NH4HS  and  in  concentrated  HC1. 

2.  KOH  or  NaOH  precipitates  white  stannic  acid — H2SnO3, 
soluble  in  excess ;  on  boiling  no  reprecipitation  takes  place — 
distinction  from  stannous  salts. 

3.  NH4OH  produces  the  same  precipitate,  insoluble  in  excess. 

4.  Heated  on  charcoal,  before  the  blowpipe,  with  Na2CO3, 
metallic  tin  is  formed,  with  the  production  of  a  white  incrust- 
ation of  the  oxide. 

REACTIONS   OF   GOLD   (Au). 
Use  a  solution  of  auric  chloride  (AuCl3). 

1.  H2S  precipitates  black  auric  sulphide — Au2S3,  insoluble 
in  HC1,  soluble  in  (NH4)2S. 

2.  H2C2O4  or  FeSO4  precipitates  metallic  gold  as  a  finely 
divided  brown  powder. 

3.  SnCl2  mixed  with  SnCl4  (prepared  by  SnCl2  and  chlo- 
rine water)  produces   a   purple-red   precipitate   or   coloration 
(Purple  of  Cassius),  consisting  of  the  mixed   oxides  of  gold 
and  tin. 

4.  Heated  on  charcoal,  before  the  blowpipe,  metallic  gold 
is  produced. 

REACTIONS   OF   PLATINUM   (Pt). 
Use  a  solution  of  platinic  chloride  (PtCl4). 

1.  H2S  causes  a  brown  precipitate  of  platinic  sulphide — 
PtS2,  insoluble  in  HC1,  soluble  in  (NH4)2S. 

2.  KC1  produces  a  yellow  crystalline  precipitate  of  potas- 
sium platinic  chloride — K2PtCl6. 

3.  NH4C1  forms  a  similar  precipitate  of  (NH4)2PtCl6.     The 


BASES. 


55 


precipitation  in  both  cases  is  facilitated  by  the   addition   of 
alcohol. 

4.  Zn,    Fe   and    some   other    metals    precipitate    metallic 
platinum. 

5.  Heating  on  charcoal,  before  the  blowpipe,  produces  the 
metal. 


SUMMARY  OF  TESTS  WITH  SOLUBLE  SALTS  OF  GROUP  V. 


As 

Sb 

Sn 

Au 

Pt 

H2S 

Yellow 
ppt. 
soluble  in 
NH4HS. 

Orange-red 
ppt. 
soluble  in 
NH^HS. 

Brown  or  Yellow 
ppt. 
soluble  in 
(NH4)2S. 

Black 
ppt. 
soluble  in 
(NH4)2S. 

Brown 
ppt. 
soluble  in 
(NH4)2S 

KOH 

No  change. 

White 
ppt. 
soluble  in 
excess. 

White 
ppt. 
soluble  in 
excess. 

No  change. 

With  excess 
of  HC1 
yellow 
ppt. 

NH4OH 

No  change. 

White 
ppt. 
insoluble  in 
excess. 

White 

insoluble  in 
excess. 

Red  ppt. 
fulminating 
gold. 

With  excess 
of  HCt 
yellow 
ppt. 

Marsh's  Test 

Black  spot 
with  metallic 
lustre,  soluble 
in  Ca(OCl)2 

Black,  sooty 
spot,  soluble 
in  (NHJHS, 
insoluble  in 
Ca(OCl)2. 

Reinsch's  Test 

White 
sublimate  of 
octahedral 
crystals. 

White 
amorphous 
sublimate. 

Fleitman's  Test 

Dark  spot 
of  metallic 
silver. 

No  change. 

DIRECTIONS  FOR  THE  DETECTION  OF  THE  BASES  IN  A  SOLUTION 
CONTAINING  SOLUBLE  SALTS  OF  GROUP  V. 

Add  HC1  and  H2S,  collect  the  precipitate,  transfer  to  a  dish,  boil  with 
concentrated  HC1 ;  filter. 


Ppt.  As,  Au,  Pt. 

If  yellow,  As  only  is  present ;  if  dark, 
digest  with  (NH4)2CO3,  filter. 


wash, 


Ppt.  Au,  Pt. 
Dissolve  in  HC1  and  HNO3. 
Divide  in  two  parts. 

Filt.  As. 
Acidify 
with  HC1, 
yellow 
ppt. 

Au. 
Add  SnCl2  purple. 
Confirm  by 
testing  original 
solution. 

Pt. 
Add  KC1,  yellow 
ppt.      Confirm  by 
testing  original 
solution. 

Filt.  Sb,  Sn. 

Dilute  with  H2O,  boil  in  a  dish  with  a  strip 
of  platinum  foil,  and  a  small  piece  of  zinc 
so  the  metals  touch.  The  platinum  will 
be  coated  with  black  Sb,  while  the  Sn 
will  be  deposited  as  a  black  sediment. 
This  black  sediment  is  dissolved  in  strong 
HC1  and  tested  with  HgCl2,  white  ppt. 
if  Sn  be  present.  Treat  the  foil  with  a 
few  drops  of  HNO3,  dissolve  in  solution 
H2C4H4O6,  add  H2S,  orange  ppt.  if  Sb 
be  present. 


56 


ANALYTICAL   CHEMISTRY. 


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BASES.  57 

PRECAUTIONS  AND  OBSERVATIONS  ON  THE  PRECEDING  CHART. 

A  great  deal  of  time  may  be  saved  by  carefully  noting  the 
color  of  the  precipitate  produced  by  H2S.  If  stannic  salts 
are  absent  and  the  precipitate  yellow,  As  only  is  present;  if 
orange,  Sb  is  present  and  As  may  be ;  when  such  an  orange 
precipitate  is  obtained  and  Sn(ic)  is  absent,  the  readiest 
method  of  separation  is  to  wash,  and  add  to  the  precipitate 
(NH4)2CO3.  As  will  be  dissolved,  and  may  be  detected  in 
the  filtrate  by  adding  HC1,  while  the  Sb  remains  on  the  filter. 
When  the  precipitate  is  dark,  Au  and  Pt  should  be  sought 
for  in  the  original  solution,  as  well  as  separated  by  the  chart 

Only  a  small  piece  of  zinc  is  necessary  to  effect  the  separa- 
tion of  Sn  and  Sb. 


GROUP  VI.— MERCURY  (1C),  BISMUTH,  COPPER, 
CADMIUM. 

REACTIONS  OF  MERCURY  as  mercuric  salt  (Hg(ic)). 
Use  a  solution  of  mercuric  chloride  (HgCl2). 

1.  H2S   or  NH4HS  produces,  when  in  small  proportion,  a 
whitish  precipitate  of  (HgS)2HgCl2;  a  further  addition  of  the 
reagent,  together  with   the   application    of  heat,   causes   the 
formation  of  a  black  precipitate  of  mercuric  sulphide — HgS, 
insoluble  in  either  HC1  or  HNO3,  but  soluble  in  a  mixture  of 
the  two. 

2.  KOH  or  NaOH  produces  a  yellow  precipitate  of  mer- 
curic oxide — HgO  ;  unless  the  reagent  be  in  excess,  a  brown 
precipitate  of  a  basic  salt  is  formed. 

2KOH  -f-  HgCl2  =  HgO  +  2KC1  +  H2O. 

3    NH4OH  precipitates  white  mercur-ammonium  chloride 
— NH2HgCl. 

2NH4OH  +  HgCl2  =  NH2HgCl  +  NH4C1  -f  2H2O. 
This  precipitate  is  readily  soluble  in  HC1  and  in  HC2H3O2. 

4.  K2CrO4  produces  a  red  precipitate  of  mercuric  chro- 
mate — HgCrO4. 

5.  KI  precipitates  mercuric  iodide — HgI2,  first  yellow,  but 
rapidly  becoming  scarlet.     This  precipitate  is  readily  soluble 
in  excess  of  KI  or  HgCl2. 

6.  SnCl2,  in  small  quantity,  in  the  presence  of  HC1,  pre- 


58  ANALYTICAL    CHEMISTRY. 

cipitates  mercurous  chloride — Hg2Cl2.  On  the  addition  of 
a  larger  quantity  of  the  reagent,  the  mercurous  chloride  is 
reduced  to  the  metal,  which  may  be  collected  into  a  globule. 

7.  Na2CO3  produces  a   reddish-brown  precipitate  of  basic 
carbonate— HgCO3(HgO)3. 

8.  Before  the  blowpipe,  HgO  breaks  up  into  Hg  and  O. 
HgS,  under  similar  circumstances,  sublimes  unchanged. 

REACTIONS  OF  BISMUTH  (Bi). 
Use  a  solution  of  bismuth  nitrate  (Bi(NO3)3). 

1.  H2S  or  NH4HS  produces  a  black  precipitate  of  bismuth 
trisulphide  —  Bi2S3,  insoluble   in    dilute   acids    and   alkalies, 
soluble  in  boiling  HNO3. 

2.  KOH,  NaOH  or  NH4OH  forms  a  white  precipitate  of 
bismuth  hydrate — Bi(OH)3,  converted  by  boiling  into  the 
yellow  oxide — Bi2O3. 

3.  K2CrO4   precipitates   yellow  bismuth    chromate  — 
Bi2(Cr04)3. 

4.  KI  forms  a  brown  precipitate  of  bismuth  iodide — BiI3, 
soluble  in  excess  of  the  reagent. 

5.  Na2CO3  precipitates  white  oxycarbonate  of  bismuth — 
(BiO)2C03.H20. 

3Na2CO3  +  2  Bi(NO3)3  -f  H2O  =  (BiO)2CO3  H2O  +  6NaNO3  +  2CO2. 

6.  H2O  in  excess,  when  there  is  not  an  excess  of  free  acid, 
precipitates  bismuth  subnitrate — BiONO3.H2O. 

Bi(NO3)3  -f  2H2O  =  BiONO3.H2O  +  2HNO3. 

When  the  chloride  is  so  diluted  the  oxychloride — BiOCl, 
separates. 

7.  Bismuth  on  charcoal,  before  the  blowpipe,  forms  a  hard 
bead  of  metal,  with  a  characteristic  incrustation  of  oxide — deep 
orange-yellow  while  hot,  pale  when  cold. 

REACTIONS  OF  COPPER  (Cu). 
Use  a  solution  of  cupric  sulphate  (CuSO4). 
i.  H2S  and  NH4HS  precipitate  black  cupric  sulphide — 
CuS,  insoluble  in  dilute  acids  and  alkalies,  slightly  soluble  in 
NH4HS,  and  entirely  dissolved  by  boiling  HNO3.     This  pre- 
cipitation is  prevented  by  KCN. 


BASES.  59 

2.  KOH   or   NaOH    produces  a  light  blue   precipitate   of 
cupric  hydrate — Cu(OH)2,  insoluble  in  excess,  and  converted 
by  boiling  into  black   cupric  oxyhydrate — (CuO)2Cu(OH)2. 
In  the  presence  of  non-volatile  organic  acids  this  precipitation 
does  not  take  place,  but  a  blue  color  results. 

2KOH  +  CuSO4  =  Cu(OH)2  -f  K2SO4. 
3Cu(OH)2  =  (CuO)2Cu(OH)2  +  2H2O. 

3.  NH4OH  in  small  quantity  forms  a  greenish-blue  precipi- 
tate, readily  soluble  in  excess,  forming  tetra-ammonio-cupric 
sulphate— (NH3)4CuSO4.H2O. 

4NH4OH  +  CuSO4  =  (NH3)4CuSO4.H2O  +  sH2O. 

4.  K4Fe(CN)6    precipitates    reddish-brown     cupric    ferro- 
cyanide — Cu2Fe(CN)6. 

5.  Metallic  Fe  or  Zn  precipitates  red  metallic  Cu. 

6.  In  the  outer  blowpipe  flame  copper  salts  color  the  borax 
bead  green  while  hot,  blue  when  cold.     In  the  inner  flame, 
after  moistening  with  SnCl2,  it  becomes  red,  owing  to  formation 
of  Cu2O. 

REACTIONS   OF  CADMIUM  (Cd). 

Use  a  solution  of  cadmium  sulphate  (CdSO4). 
# 

1.  H2S  or  NH4HS  precipitates  yellow  cadmium  sulphide 

— CdS,  soluble  in  hot   HNO3,  but   insoluble  in   NH4HS  or 
KCN. 

2.  KOH  or  NaOH  produces  a  white  precipitate  of  cadmium 
hydrate — Cd(OH)2,  insoluble  in  excess. 

3.  NH4OH  causes  the  same  white  precipitate  of  Cd(OH)2, 
soluble  in  excess. 

4.  Na2CO3  produces  a  white  precipitate  of  cadmium  car- 
bonate— CdCO3,  insoluble  in  excess,  but  slightly  soluble  in 
ammonium  salts,  entirely  soluble  in  NH4OH. 

5.  On  charcoal,  before  the  blowpipe,  the  salts  of  cadmium 
are  reduced  to  metal  and  volatilize,  forming  a  brownish  in- 
crustation of  oxide. 


60  ANALYTICAL   CHEMISTRY. 

SUMMARY  OF  TESTS  WITH  SOLUBLE  SALTS  OF  GROUP  VI. 


Hg(ic) 

Bi 

Cu 

Cd 

H2S  or  NH4HS 

Black  ppt. 
insoluble  in  HNO3. 

Black  ppt. 

Black  ppt. 
soluble  in  KCN. 

Yellow  ppt. 
insoluble  in  KCN. 

KOH 

Yellow  ppt. 

White  ppt. 

Blue  ppt. 

White  ppt. 

NH4OH 

White  ppt. 

White  ppt. 

Blue  ppt. 
soluble  in  excess. 

White  ppt. 
soluble  in  excess. 

Na2C03 

Reddish-brown  ppt. 

White  ppt. 

Blue  ppt. 

White  ppt. 

DIRECTIONS  FOR  THE  DETECTION  OF  THE  BASES  IN  A  SOLUTION 

CONTAINING  SOLUBLE  SALTS  OF  GROUP  VI. 

Add  HC1  and  H2S,  collect,  boil  with  HNO3  ;  filter. 


Ppt.  Hg. 
Black. 

Fill.  Bi,  Cu,  Cd. 
Add  NH4OH  in  excess  ;  filter. 

Ppt.  Bi. 
White. 

Fill.  Cu,  Cd. 
Add  H2S,  collect  and  wash  ;  boil  with  dilute 
HaSO4;  filter. 

Ppt.  Cu. 
Black. 

Filt.  Cd. 
Dilute,  add  H2S,  yellow  ppt. 

BASES. 


61 


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Z 

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Wash,  dis 


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until  odor  of  HCN 
removed,  add  KOH 
excess,  filter. 


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Yellow  if  Cr  be 
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Collect  and 
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dilute  H2SO4, 
filter. 


62  ANALYTICAL   CHEMISTRY. 

PRECAUTIONS  AND  OBSERVATIONS  ON  THE  PRECEDING  CHART. 

1.  Unless  excess  of  H2S  be  used,  and  the  solution  warmed, 
Hg  will  not  be  thoroughly  precipitated. 

2.  When  Sn,  Au  and  Pt  are  present,  the  yellow  ammonium 
sulphide  (NH4)2S  must  be  used  to  dissolve  Group  V.     When 
Cu   is  present  some  of  it  may  be  dissolved,  so  that  any  dark 
precipitate  in  that  group  should  be  tested  for  Cu. 

3.  When  the  precipitate  insoluble  in  NH4HS  is  boiled  with 
HNO3,    a   black    mass    of    sulphur    is    sometimes    obtained. 
This   is    readily   distinguished   from    Hg,   because    it    floats 
as    one    mass    on   the    liquid,  while  the  mercury  collects  at 
the  bottom  as  a  heavy,  black  precipitate. 


GROUP   VII.— SILVER,   MERCURY(OUS),   LEAD. 

REACTIONS  OF  SILVER  (Ag). 
Use  a  solution  of  silver  nitrate  (AgNO3). 

1.  HC1  or  soluble  chlorides  precipitate  white,  curdy  silver 
chloride — AgCl,  insoluble  in  HNO3,  soluble  in  NH4OH,  form- 
ing ammonio-silver  chloride — (AgCl)2(NH3)3,  from  which 
the  chloride  is  again  precipitated  by  acids. 

HC1  +  AgNO3  =  AgCl  -f  HNO3. 
2AgCl  +  3NH4OH  =  (AgCl)2(NH3)3  +  3H2O. 

2.  H2S  or  NH4HS  produces  a  black  precipitate  of  silver 
sulphide — Ag2S,  insoluble   in    dilute   acids  and  in    alkalies, 
soluble  in  boiling  HNO3. 

3.  KOH   or   NaOH   forms  a  grayish-brown  precipitate  of 
silver   oxide  —  Ag2O,    insoluble    in    excess,   but   soluble    in 
NH4OH. 

2KOH  +  2AgNO3  =  Ag2O  -f  2KNO3  -f  H2O. 

4.  NH4OH    in    small    quantity,  precipitates    silver  oxide, 
soluble  in  excess. 

5.  K,CrO4  produces  a  red  precipitate  of  silver  chromate — 
Ag2CrO4,  soluble  in  concentrated  HNO3  and  in  NH4OH. 

6.  KI  and  KBr  produce  precipitates  of  silver  iodide  — Agl, 
yellow,  insoluble   in  NH4OH,  and   silver  bromide  —  AgBr, 
yellowish-white,  slowly  soluble  in  NH4OH. 

7.  KCN  precipitates  white  silver  cyanide — AgCN,  soluble 
in  excess  and  in  concentrated  HNOS. 


BASES.  63 

8.  Heated  with  Na2CO3  'on  charcoal,  before  the  blowpipe, 
compounds  of  silver  form  a  bright,  metallic  button,  soluble  in 
HN03. 

REACTIONS  OF  MERCURY  as  mercurous  salt  (Hg(ous)). 
Use  a  solution  of  mercurous  nitrate  (Hg2(NO3^2). 

1.  HC1  or  soluble  chlorides  precipitate  mercurous  chloride 
— Hg2Cl2,  converted  by  strong  HNO3  into  a  mixture  of  HgCl2 
and  Hg(NO3)2,  also  becoming  black  on  the  addition  of  NH4OH, 
forming  NH2Hg2CL 

2.  H2S  or  NH4HS  precipitates  a  mixture  of  Hg  with  HgS. 

3.  KOH  or  NaOH  produces  a  black  precipitate  of  mer- 
curous oxide — Hg2O,  insoluble  in  excess. 

4.  NH4OH    causes    a    black    precipitate    of    mercurous- 
ammonium  nitrate — NH2Hg2NO3. 

2NH4OH  +  Hg2(N03)2  =  NH2Hg2N03  +  NH4NO3  -f  2H2O. 

5.  K2CrO4    forms   an    orange    precipitate    of    mercurous 
chromate — Hg2CrO4, 

6.  KI  precipitates  green  mercurous  iodide — Hg2I2. 

7.  Before   the   blowpipe,  mercurous    salts   volatilize,  some 
being   converted    into    mercuric    salt   and    mercury,  both   of 
which  sublime. 

REACTIONS   OF   LEAD  (Pb). 
Use  a  solution  of  lead  acetate  (Pb(C2H3O2)2). 

1.  HC1,  or  soluble  chlorides,  produce  a  white  precipitate  of 
lead  chloride — PbCl2,  soluble  in  hot  water. 

2.  H2S  or  NH4HS  precipitates  black  lead  sulphide — PbS, 
insoluble  in  HC1,  soluble  in  hot  HNO3. 

3.  KOH  or  NaOH  produces  a  white  precipitate  of  lead 
hydrate — Pb(OH)2,  soluble  in  large  excess,  forming  potas- 
sium or  sodium  plumbate — K2PbO2  or  Na2PbO2. 

4.  NH4OH  precipitates  white  basic  lead  hydrate. 

5.  K2CrO4  produces  a  yellow  precipitate  of  lead  chromate 
-PbCrO4,  soluble  in  KOH  and  in  strong  HNO3. 

6.  KI    forms   a   yellow  precipitate  of  lead   iodide — PbI2, 
soluble  in  boiling  water. 

7.  H2SO4  produces  a  white  precipitate  of  lead  sulphate — 


64 


ANALYTICAL   CHEMISTRY. 


PbSO4,  insoluble  in  acids,  but  soluble  in  solution  of  ammo- 
nium acetate  or  tartrate. 

8.  Na2CO3    precipitates    white    basic    lead    carbonate  — 
(PbC03)2Pb(OH)2. 

3Na2C03  +  3Pb(C2H302)2  +  H20  =  (PbCO3)2Pb(OH)2 
+  6NaC2H3O2  -f  CO2. 

9.  Before  the  blowpipe,  on  charcoal,  lead  compounds  are 
converted    into  a  malleable    globule  of  the  metal,  with  the 
formation  of  some  yellow  oxide. 


SUMMARY  OF  TESTS  WITH    SOLUBLE  SALTS  OF  GROUP  VII. 


Ag 

Hg(ous) 

Pb 

HC1 

White  ppt. 
soluble  in  NH4OH. 

White  ppt.,  turning 
black  with  NH4OH. 

White  ppt. 
soluble  in  hot  H2O. 

H2S  and  NH4HS 

Black  ppt. 

Black  ppt. 

Black  ppt. 

KOH 

Brown  ppt. 

Black  ppt. 

White  ppt. 

NH4OH 

Brown  ppt. 

Black  ppt. 

White  ppt. 

Na2CO3 

Brown  ppt. 

Black  ppt. 

White  ppt. 

K2CrO4 

Red  ppt. 

Orange  ppt. 

Yellow  ppt. 

KI 

Yellow  ppt. 

Green  ppt. 

Yellow  ppt. 

DIRECTIONS   FOR  THE  DETECTION  OF  THE  BASES   IN  A  SOLUTION 
CONTAINING  SOLUBLE  SALTS  OF  GROUP  VII. 
Add  HC1,  collect,  wash  and  pour  on  the  filter  boiling  H2O. 


Ppt.  Ag,  Hg(ous). 
Pour  on  the  filter  NH4OH. 

Filt.  Pb. 
Add  H2SO4,  white  ppt. 

Ppt.  Hg(ous). 
Black. 

Filt.  Ag. 
Add  HNO3  in  excess, 
white  ppt. 

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66  ANALYTICAL   CHEMISTRY. 

SPECIAL   PRECAUTIONS   TO   BE   OBSERVED    IN   THE   EXAMI- 
NATION  FOR   BASES. 

1.  HC1  may  precipitate  basic  salts  of  Bi  and  Sb;  these  are 
readily  distinguished  from  the  bases  of  Group  VII  by  dis- 
solving on  the  further  addition  of  HC1. 

2.  The  precipitation  with  HC1  should  be  performed  in  the 
cold,  in  order  to  prevent  the  PbCl2  from  dissolving. 

3.  Before  commencing  the  analysis  of  a  solution  its  action 
on  litmus  paper  should  always  be  noted.     If  alkaline  a  pre- 
cipitate may  be  caused  by  HC1,  on  account  of  the  presence  of 
substances   soluble  in  alkalies,  as  As2S3  in  (NH4)2CO3,  and 
AgCl  in  NH4OH.     Or  silicic  acid  from  alkaline  silicates  may 
be  precipitated,  in  which  case  it  is  necessary  to  precede  the 
addition  of  HC1,  by  evaporation  to  dryness  with  HNO3,  filter- 
ing out  the  insoluble  silica,  and  proceeding  with  the  filtrate  in 
the  usual  way. 

4.  The  addition  of  H2S    should  be  slow  and   interrupted 
from  time  to  time,  to  warm  and  agitate  the  solution,  and  so 
continued  until  it  smells  strongly  of  the  reagent. 

5.  When  the  presence  of  HNO3  is  suspected,  the  filtrate 
after  the  precipitation  of  Group  VII,  and  before  the  addition 
of  H2S,  should  be  evaporated  to  dryness  to  expel  the  HNO3, 
dissolved  in  water  and,  if  necessary,  acidified  with  HCL 

SPECIAL  OBSERVATION. 

The  separation  into  groups  is  accomplished  by  the  following 
reagents,  which  are  called  group  reagents  :  — 


HCl  precipitates  Ag,  Hg(ous),  Pb,    ...........  Group  VII. 

f  Insoluble  in  (NH4)2S,  Hg(ic),  Bi,  Cu,  Cd,    .  Group  VI. 
2  I  Soluble  in  (NH4)aS,  As,  Sb,  Sn,  Au,  Pt,  .    .  Group  V. 

f  Hydrates  insoluble  in  NH4OH,  Fe,  Ce,  Al,  Cr,  Group  IV. 
H±H  I  Hydrates  soluble  in  NH4OH,  Mn,  Zn,  Co,  Ni,  Group  III. 

<NH^°»       "       Ba,Sr,Ca>l 
Na2HPO4         "       Mg,  j 

Notprecipitatedina  |  R>  N      ^  _   G         L 

group 

When  one  of  these  reagents  fails  to  produce  a  precipitate,  it 
indicates  the  absence  of  that  group,  and  the  next  group 
reagent  is  immediately  added.  , 


ACIDS.  67 


SECTION  II. 

ACIDS. 

GROUP  I.— HYDROCHLORIC  ACID,  HYDROBRO- 
MIC  ACID,  HYDRIODIC  ACID,  HYDROCY- 
ANIC ACID,  HYDROFLUORIC  ACID. 

REACTIONS   OF   HYDROCHLORIC   ACID  (HC1). 
Use  a  solution  of  potassium  chloride  (KC1). 

1.  AgNO3  produces  a  white,  curdy  precipitate  of  silver 
chloride  —  AgCl,   insoluble    in    HNO3,   readily   soluble    in 
NH4OH. 

This  precipitate  should  be  preserved  to  compare  with  those 
of  AgBr  and  Agl. 

2.  Hg2(NO3)2   causes    a    white    precipitate   of  mercurous 
chloride  —  Hg2Cl2,   insoluble   in    HNO3,  blackening   on   the 
addition  of  NH4OH. 

3.  Pb(C2H3O2)2  forms  a  white  crystalline  precipitate  of  lead 
chloride — PbCl2,  soluble  in  33  parts  of  boiling  water. 

4.  Warming  with  H2SO4  and  MnO2  causes  the  evolution  of 
Chlorine,  recognized  by  its  odor  and  color. 

5.  On  warming  with  H2SO4,  hydrochloric  acid  is  given  off, 
recognized  by  its  odor  and  intensely  acid  reaction;  also  by 
the  dense  white  fumes  of  NH4C1  produced  by  holding  a  rod 
moistened  with  NH4OH  near  the  mouth  of  the  tube. 

REACTIONS   OF   HYDROBROMIC   ACID  (HBr). 
Use  a  solution  of  potassium  bromide  (KBr). 

1.  AgNO3  produces  a  yellowish-white  precipitate  of  silver 
bromide  —  AgBr,   insoluble    in    HNO3,   slowly   soluble    in 
NH4OH. 

This  precipitate  should  be  compared  with  those  of  AgCl  and 
Agl. 

2.  HgCl2  forms  a  white  precipitate  of  mercuric  bromide — 
HgBr2,  soluble  in  a  large  quantity  of  water. 

3.  Hg2(NO3)2  precipitates  yellowish-white  mercurous  bro- 
mide— Hg2Br2. 


68  ANALYTICAL    CHEMISTRY. 

4.  H2SO4,  when  concentrated,  causes  the  evolution  of  red 
vapors  of  bromine ;  this  occurs  more  readily  in  the  presence 
of  MnO2. 

5.  Chlorine  water  and  starch  paste  cause  a  yellow  color, 
due  to  formation  of  starch  bromide ;  in  dilute  solutions  it  is 
necessary  to  agitate  the  mixture  with   ether   or   chloroform, 
which  will  separate,  carrying  the  bromine  in  solution,  with  a 
red  or  reddish-brown  color. 

REACTIONS   OF   HYDRIODIC   ACID  (HI). 
Use  a  solution  of  potassium  iodide  (KI). 

1.  AgNO3  produces  a  yellowish  precipitate  of  silver  iodide 
— Agl,  insoluble  in  HNO3  and  almost  insoluble  in  NH4OH. 

2.  Hg2(NO3)2  precipitates  green  mercurous  iodide — Hg2I2. 

3.  HgCl2  causes  a  red  precipitate  of  mercuric  iodide — HgI2, 
soluble  in  excess  of  either  reagent. 

4.  Pb(C2H3O2)2  produces  a  yellow  precipitate  of  lead  iodide 
— PbI2,  soluble  in  boiling  water. 

5.  Chlorine  water  and  starch  paste  form  a  blue  color  of 
starch  iodide,  which  color  disappears  on  heating  and  returns 
on  cooling,  also  destroyed  by  excess  of  chlorine  water. 

6.  A  concentrated  solution  of  one  part  CuSO4  and  three 
parts  FeSO4  produces  a  grayish  precipitate  of  cuprous  iodide 
— Cu2I2.     This   reaction  is   useful   in   separating  iodine   from 
chlorine  and  bromine. 

HYDROFLUORIC   ACID  (HF). 

The   evolution  of  intensely  irritating  fumes  of  this  acid,  on  the  addition  of 
H2SO4  to  calcium  fluoride,  which  etch  glass,  is  sufficiently  characteristic. 

REACTIONS   OF   HYDROCYANIC   ACID  (HCN). 
Use  a  solution  of  potassium  cyanide  (KCN). 

1.  AgNO3  produces  a  white  precipitate  of  AgCN,  soluble 
in  KCN,  sparingly  soluble  in  NH4OH,  and  insoluble  in  dilute 
HN03. 

2.  NH4HS  evaporated  with  KCN  to  dryness,  on  a  water 
bath,  will  give  on  dissolving  in  water  and  adding  Fe2Cl6  a 
deep  blood  red  color,  due  to  formation  of  sulphocyanate. 

3.  Fe2Cl6  and  FeSO4,  then  NaOH  until  a  precipitate  is  pro- 


ACIDS. 


69 


duced,  heat  and  add  HC1,  will  give  a  deep  blue  residue  of 
ferric  ferrocyanide  (Prussian  blue). 

SUMMARY  OF  TESTS  WITH  SOLUBLE  SALTS  OF  GROUP  I. 


HC1 

HBr 

HI 

HCN 

AgN03 

White  ppt. 
soluble  in  NH4OH 

Yellowish-  white 
ppt.  slowly  soluble 
in  NH4OH. 

Yellowish  ppt. 
insoluble  in 
NH4OH. 

White  ppt. 
sparingly  soluble  - 
in  NH4OH. 

Hg2(N03)3 

White  ppt. 

Yellowish-white 
ppt. 

Greenish  ppt. 

White  ppt. 

HgCla 

No  ppt. 

White  ppt. 

Red  ppt. 

No  ppt. 

Pb(C2H302)2 

White  ppt. 

White  ppt. 

Yellow  ppt. 

White  ppt. 

Chlorine  water 
and 
starch  solution. 

No  color. 

Yellow  color. 

Blue  color. 

No  color. 

DIRECTIONS  FOR  THE  DETECTION  OF  THE  ACIDS  IN  A  SOLU- 
TION CONTAINING  SOLUBLE  SALTS  OF  GROUP  I. 

Add  HNO3  and  boil  for  some  time;  the  HCN  is  driven  off. 
This  operation  should  be  performed  with  caution,  so  as  to  avoid 
inhaling  these  poisonous  fumes. 

After  all  the  HCN  has  been  driven  off,  add  AgNO3  and 
shake  well ;  collect  the  precipitate  of  AgCl,  AgBr,  Agl  on 
a  filter,  and,  after  washing,  pour  on  the  filter  NH4OH ;  AgCl 
will  be  dissolved,  and  may  be  detected  in  the  filtrate  by 
acidifying  with  HNO3. 

To  detect  HBr  and  HI  take  another  portion  of  the  original 
solution ;  add  a  few  drops  of  starch  solution,  and  then,  slowly, 
chlorine  water,  until,  after  agitation,  the  liquid  smells  of  it. 
If  there  is  a  blue  color  HI  is  present.  Continue  the  addition 
of  chlorine  water  until  the  blue  color  is  destroyed,  when,  if 
the  solution  is  yellow,  HBr  is  indicated.  This  may  be  still 
further  verified  by  agitating  with  chloroform,  which  will,  after 
separating,  assume  a  red  or  red  brown  color  if  HBr  is  present. 


70  ANALYTICAL    CHEMISTRY. 

GROUP   II. 


HYPOCHLOROUS  ACID. 
CHLORIC  " 

HYDROXYL. 

HYDROSULPHURIC  ACID. 
SULPHUROUS 
SULPHURIC 
THIOSULPHURIC 
NITRIC 


HYPOPHOSPHOROUS  ACID, 

ORTHOPHOSPHORIC 

PYROPHOSPHORIC  " 

METAPHOSPHORIC 

BORIC 

CARBONIC 

SILICIC 


HYPOCHLOROUS  ACID  (HC1O). 

This  acid  is  known  in  combination  with  calcium  Ca(ClO)2,  bleaching  powder, 
and  with  sodium  NaCIO,  Labarraque's  solution.  The  chlorine-like  odor  devel- 
oped on  the  addition  of  an  acid  is  sufficient  evidence  of  its  presence. 

CHLORIC   ACID   (HC1O3). 

The  compounds  of  potassium  and  sodium  with  this  acid  are  best  known. 

1.  H2SO4,  added  to  one   of  them,  causes  the  evolution  of  yellow  chlorine 
tetroxide — C12O4,  having  a  characteristic  odor.     When  the  dry  salt  is  used  this 
reaction  takes  place  with  explosive  violence. 

2.  AgNO3   produces  no   precipitate.     This   is   an  important  distinction  from 
HC1. 

WATER  (H2O),   HYDRATES    AND    OXIDES. 

Water  is  distinguished  by  having  no  odor  or  taste,  not 
changing  litmus,  and  evaporating  without  residue. 

The  soluble  hydrates  KOH,  NaOH,  LiOH,  Ca(OH)2, 
Sr(OH)2,  Ba(OH)2,  yield  solutions  with  water  which  change 
red  litmus  paper  blue.  The  insoluble  hydrates  give  off  steam 
when  heated  in  a  dry  tube. 

The  soluble  oxides  K2O,  Na2O,  Li2O,  CaO,  SrO,  BaO,  are 
known  by  forming  hydrates  with  water.  The  insoluble  oxides 
are  recognized  by  giving  negative  tests  for  acids  when  dis- 
solved and  tested  in  the  usual  way. 

HYDROSULPHURIC   ACID  (H2S). 

Sufficient  evidence  of  the  presence  of  this  acid  is  afforded 
by  the  characteristic  odor.  In  the  case  of  sulphides,  first 
adding  H2SO4  and  warming  if  necessary.  A  trace  may  be 
detected  by  holding  over  the  mouth  of  the  tube  a  piece  of 
filter  paper  moistened  with  lead  acetate,  which  will  become 
black  (lead  sulphide)  in  the  presence  of  H2S. 


ACIDS.  71 

REACTIONS   OF   SULPHUROUS   ACID  (H2SO3). 

1.  In  solution,  uncombined,  this  acid  is  recognized  by  its 
odor   of  burning   sulphur,   by   strong   bleaching   action,   by 
decolorizing  potassium  permanganate  solution,  and  by  causing 
the  evolution  of  hydrogen  sulphide  when  added  to  a  mixture 
of  zinc  and  hydrochloric  acid. 

2.  Sulphites  are  distinguished  by  the  characteristic  odor  of 
SO2  on  the  addition  of  a  strong  acid. 

3.  Salts   of  Ag,    Hg  or    Pb,   produce    precipitates   which 
blacken  on  heating,  owing  to  formation  of  sulphides. 

4.  BaCI2  with  neutral  solutions  forms  a  white  precipitate  of 
barium  sulphite — BaSO3,  soluble  in  HC1. 

REACTIONS   OF  SULPHURIC   ACID  (H2SO4). 
Use  dilute  H2SO4,  or  an  alkali  sulphate. 

1.  BaCl2  produces  a  white  precipitate  of  barium  sulphate 
— BaSO4,  insoluble  in  boiling  concentrated  acids. 

2.  Pb(C2H3O2)2  causes  a  precipitate  of  white  lead  sulphate, 
insoluble  in  dilute  acids,  but  soluble  in  hot  concentrated  acids. 
Alcohol  increases  the  delicacy  of  this  reaction. 

3.  A  sulphate  fused  on  charcoal  with   Na2COa,  the  fused 
mass  placed  on  a  bright  silver  coin  and  moistened  with  a  drop 
of  dilute  HC1  will  cause  a  black  stain,  due  to  formation  of 
silver  sulphide.     This  reaction  is  especially  adapted  to  the 
detection  of  insoluble  sulphates. 

REACTIONS   OF  THIOSULPHURIC  (HYPOSUL- 

PHUROUS)  ACID  (H2S2O3). 
Use  a  solution  of  sodium  thitfsulphate  (Na^Os). 

1.  H2SO4  causes  the  evolution  of  sulphurous  oxide — SO2, 
recognized  by  the.  odor.     A  deposit  of  sulphur  takes  place 
at  the  same  time,  which    is   an    important  distinction    from 
sulphites. 

2.  AgNO3   produces   a   white   precipitate  of  silver  thio- 
sulphate — Ag2S2O3,  soluble  in  excess.     After  a  time  (immedi- 
ately on  heating)  the  precipitate  becomes  dark  and  then  black, 
silver  sulphide  and  sulphuric  acid  being  formed. 

3.  BaCl2  produces  a  white  precipitate  soluble  in  excess  of 
H2O,  and  decomposed  by  HC1. 


72  ANALYTICAL   CHEMISTRY. 

4.  Added   to   a   mixture   of  zinc   and   hydrochloric   acid, 
hydrogen  sulphide  is  evolved. 

REACTIONS   OF   NITRIC   ACID  (HNO3). 
Use  a  solution  of  potassium  nitrate  (KNO8). 

1.  H2SO4,  on  heating,  will  cause  the  nitric  acid  to  volatilize. 
If  copper  turnings  be  added  with  the  sulphuric  acid  colorless 
nitrogen  dioxide — N2O2  will  be  given  off,  which  in  contact 
with   air   will   form    red   nitrogen  tetroxide — N2O4,  readily 
recognized   by   the    color   and    odor.     If   alcohol   be   added 
to   the   mixture   the   characteristic  odor  of  nitrous  ether  is 
developed. 

2.  FeSO4,  acidified  with  H2SO4,  added  in  a  test  tube,  so  as 
to  form  a  layer  on  a  solution  of  a  nitrate,  acidified  with  H2SO4, 
will  cause  a  dark  layer  to  form  at  the  line  of  contact. 

3.  Indigo  solution,  strongly  acidified  with  H2SO4,  is  decolor- 
ized by  a  nitrate. 

4.  Zinc  and  potassium  hydrate  cause  the  reduction  of  the 
acid  radical  to  ammonia,  which  may  be  detected  in  the  usual 
way.     This  is  a  valuable  test  for  distinguishing  nitric  acid  in 
the  presence  of  chloric  acid. 

5.  Heated  on  charcoal,  deflagration  takes  place,  the  char- 
coal burning  at  the  expense  of  the  oxygen  of  the  nitrate. 

HYPOPHOSPHOROUS  ACID  (HH2PO2). 

1.  By  ignition  the  hypophosphites  are  resolved  into  spontaneously  inflammable 
hydrogen  phosphide — PH3,  and  phosphate. 

2.  AgNO3  produces  at  first  a  white  precipitate  of  silver  hypophosphite — 
AgH2PO2,  which  soon  becomes  black,  owing  to  formation  of  metallic  silver. 

3.  HgQ2  in  excess  causes  a  white  precipitate  of  mercurous  chloride  Hg2Cl2> 
this  takes  place  more  rapidly  on  warming,  and  in  the  presence  of  HC1. 

REACTIONS  OF  ORTHOPHOSPHORIC  ACID  (H3PO4). 
Use  a  solution  of  sodium  phosphate  (Na2HPO4). 

1.  AgNO3  causes  a  light  yellow  precipitate  of  silver  phos- 
phate—Ag3PO4,  soluble  in  HNO3  and  in  NH4OH. 

2.  Fe2Cl6,  in  presence  of  sodium   acetate,  produces  a  yel- 
lowish-white,   gelatinous    precipitate  of  ferric   phosphate — 
Fe2(PO4)2.     An  excess  of  Fe2Cl6  must  be  avoided. 


ACIDS.  73 

3.  (NH4)2MoO4,  in  neutral  or  acid  solution,  causes  a  yellow 
precipitate  to  separate  slowly,  which  is  ammonium  phospho- 
molybdate— (NH4)3PO4(MoO3)10  +  2H2O,  insoluble  in  HNO3, 
soluble  in  NH4OH. 

4.  Magnesia  mixture  (consisting  of  MgSO4,  NH4C1,  NH4OH) 
causes  a  white  precipitate  of  ammonium  magnesium  phos- 
phate—MgNH4PO4  +  6H2O.     Agitation   facilitates   the   for- 
mation of  this  precipitate. 

5.  BaCl2,  in  neutral  solution,  produces  a  white  precipitate  of 
barium  hydrogen  phosphate — BaHPO4. 

6.  Albumen  (white  of  egg)  does  not  cause  a  precipitate. 

REACTIONS  OF  PYROPHOSPHORIC  ACID  (H4P2O7). 
Use  a  solution  of  sodium  pyrophosphate  (Na4P2O7). 

1.  AgNO3  precipitates  white  silver  pyrophosphate — Ag4- 
P2O7,  soluble  in  HNO3  and  NH4OH. 

2.  MgSO4  causes  a  white  precipitate  of  magnesium  pyro- 
phosphate— Mg2P2O7,  soluble  in  excess  of  the  reagent ;  from 
this  solution   NH4OH  does  not  reprecipitate  it  in  the  cold. 
This  reaction  may  be  used  to  distinguish  ortho-  from  pyro- 
phosphoric  acid. 

3.  Neither  (NH4)2MoO4  nor  albumen  produces  a  precipitate. 

4.  BaCl2,  in    neutral    solution,  precipitates  white    barium 
pyrophosphate — Ba2P2O7. 

REACTIONS   OF   METAPHOSPHORIC  ACID  (HPOS). 
Use  a  solution  of  sodium  metaphosphate  (NaPO3). 

1.  AgNO3  produces  a  white  precipitate  of  silver  meta- 
phosphate— AgPO3,  soluble  in  HNO3  and  NH4OH. 

2.  Albumen  forms  a  white  precipitate  with  the  free  acid, 
and  with  the  salts  on  the  addition  of  acetic  acid.     This  is  an 
important  distinction  from  the  ortho-  and  pyro-acids. 

3.  Neither  (NH4)2MoO4  nor  magnesia  mixture  produces  a 
precipitate;    should  one  form  with  the  latter  reagent,   it  is 
readily  soluble  in  NH4C1. 

4.  BaCl2  forms  a  white  precipitate  of  barium  metaphos- 
phate—Ba(PO3)2. 


74  ANALYTICAL   CHEMISTRY. 

5.  Solutions  of  the  meta-  and  pyro-acids  in  water  are  con- 
verted into  the  ortho-acid  by  boiling. 

REACTIONS   OF   BORIC   ACID   (H3BO3). 
Use  a  solution  of  sodium  borate  (Na2B4O7). 

1.  AgNO3  produces  a  white  precipitate  of  silver  borate, 
soluble  in  HNO3. 

2.  BaCl2  precipitates  white  barium  borate,  soluble  in  excess 
of  water  and  in  NH4C1. 

3.  H2SO4  or  HC1  causes  the  separation  of  the  acid,  H3BO3, 
in  crystalline  form,  from  strong  solutions. 

4.  Alcohol  added  to  the  acid,  and  ignited,  burns  with  a 
characteristic  green  flame.     In  the  case  of  salts,  the  addition 
of  alcohol  is  preceded  by  that  of  H2SO4,  in  order  to  liberate 
the  free  acid.     Salts  of  copper  when  present  should  be  re- 
moved before  this  test  is  applied,  as  they  likewise  impart  a 
green  color  to  the  alcohol  flame. 

REACTIONS   OF   CARBONIC  ACID    (H2CO3). 
Use  a  solution  of  sodium  carbonate  (Na2CO3). 

1.  All  free  acids  except  HCN  and  H2S  decompose  carbon- 
ates with  effervescence.     The  escaping  gas  passed  into  a  solu- 
tion of  Ba(OH)2  or  Ca(OH)2  causes  a  white  precipitate. 

2.  BaCl2,  in    neutral    solution,  precipitates  .white   barium 
carbonate — BaCO3,  soluble  in  acids. 

3.  CaCl2   precipitates   white   calcium   carbonate — CaCO3, 
soluble  in  acids  with  effervescence. 

REACTIONS   OF   SILICIC   ACID  (H4SiO4). 

1.  Insoluble  silicates  are  determined  by  fusing  on  platinum 
foil  some  of  the  fine  powder  with  Na2CO3,  treating  the  fused 
mass  with  H2O  and  HC1,  evaporating  to  dryness  and  redis- 
solving  in  H2O,  when  the  silica,  SiO2,  will  remain  as  a  fine 
white  precipitate. 

2.  Before  the  blowpipe,  with  a  bead  of  microcosmic  salt, 
silica  forms  the  so-called  silica  skeleton,  which  is  very  charac- 
teristic. 


ACIDS.  75 

3.  Soluble  silicates  give  gelatinous  precipitates   of  silicic 
acid — H4SiO4,  on  the  addition  of  H2SO4  or  HC1;  on  evapor- 
ating this  to  dryness  with  a  little  HC1,  and  redissolving  in 
H2O,  silica — SiO2,  remains. 

4.  NH4C1  also  precipitates  H4SiO4  when  added  to  a  soluble 
silicate. 

REACTIONS   OF    HYDROFERROCYANIC   AND 

HYDROFERRICYANIC   ACIDS. 
These  acids  have  been  sufficiently  characterized  under  iron. 

DIRECTIONS  FOR  THE  DETECTION  OF  THE  ACIDS  IN  A  SOLU- 
TION  CONTAINING  SOLUBLE   SALTS  OF  GROUP   II. 

I.  Try  the  solution  with  litmus  paper ;  if  alkaline,  hydrates, 
carbonates,  borates,  silicates  and,  possibly,  phosphates,  may  be 
present.     If  the  solution  is  acid,  neutralize  with  NaOH  before 
applying  the  following  tests  : — 

II.  Evaporate  a  portion  of  the  solution  to  dryness,  and  add 
concentrated  H2SO4.     The  following  acids  will  give  character- 
istic reactions,  and  may  be  recognized  by  the  further  applica- 
tion of  tests  previously  given : — 

HC1O  and  HC1O3  give  the  odor  of  chlorine.  They  are 
readily  distinguished  by  the  chlorate  deflagrating  with  char- 
coal, by  its  liberating  oxygen  on  heating,  and  by  the  yellowish- 
green  gas,  C12O4,  which  is  given  off.  The  addition  of  H2SO4 
to  a  chlorate  should  be  performed  with  very  small  quantities 
and  with  great  care,  on  account  of  the  tendency  of  the  chlorine 
tetroxide  to  decompose  with  explosive  violence.  The  most 
characteristic  difference  of  HC1O  is  the  white  precipitate  of 
AgCl  which  it  gives  with  AgNO3. 

H2S  is  recognized  by  its  peculiar  odor,  and  by  blackening  a 
piece  of  filter  paper  moistened  with  solution  of  lead  acetate. 

H2SO3  and  H2S2O3  give  off  SO2,  readily  recognized  by  its 
odor  and  bleaching  property.  They  are  distinguished  from 
each  other  by  H2SO4  producing  with  a  concentrated  solution 
of  thiosulphate  a  white  precipitate  of  sulphur  in  addition  to 
the  liberation  of  SO2. 

HNO3  is  detected  by  the  peculiar  acid  fumes,  which  become 
red  when  metallic  copper  is  added  with  the  H2SO4.  All  the 


76  ANALYTICAL   CHEMISTRY. 

special  tests,  previously  given,  of  this  acid  should  be  applied 
before  a  conclusion  is  reached  concerning  the  presence  or 
absence  of  it. 

HH2PO2  is  readily  detected  by  its  odor. 

H3BO3  is  easily  detected  when,  in  addition  to  the  H2SO4, 
some  alcohol  is  added  and  ignited.  In  the  absence  of  salts  of 
copper,  the  green  flame  is  evidence  of  this  acid. 

H2CO3  is  recognized  by  effervescence  with  the  dilute  acid  in 
the  cold  without  odor.  When  other  gases  are  given  off  at  the 
same  time,  the  CO2  may  be  detected  by  passing  into  lime 
water,  which  will  cause  a  precipitate  of  CaCO3. 

H4SiO4  is  precipitated  as  a  white,  gelatinous  mass  when  the 
H2SO4  is  added  to  a  soluble  silicate.  Insoluble  silicates  will 
be  treated  of  later. 

H4Fe(CN)6  and  H6Fe2(CN)12  give  the  odor  of  HCN.  Their 
presence  is  confirmed  by  the  use  of  FeSO4  and  Fe2Cl6. 

III.  The  acids  of  this  group  not  detected  by  H2SO4  are  sul- 
phuric and  the  three  phosphoric  acids. 

H2SO4  is  precipitated,  on  the  addition  of  BaCl2,  as  BaSO4, 
insoluble  in  HC1  or  HNO3. 

H3PO4  is  precipitated  yellow  on  the  addition  of  AgNO3,  the 
precipitate  being  soluble  in  HNO3  (distinction  from  HC1,  HBr 
and  HI).  (NH4)2MoO4  is  the  most  distinctive  test.  It  causes 
a  yellow  precipitate  when  a  few  drops  of  the  solution  are  added 
to  some  of  the  reagent  in  a  test  tube  and  warmed  gently  (dis- 
tinction from  H4P2O7  and  HPO3).  Magnesia  mixture  is  also  a 
characteristic  test  for  this  acid. 

H4P2O7  gives  a  white  precipitate  with  AgNO3.  It  is  further 
distinguished  from  the  ortho-acid  by  (NH4)2MoO4,  and  from 
the  meta-acid  by  its  behavior  with  MgSO4,  and  with  albumen. 
HPO3  also  gives  a  white  precipitate  with  AgNO3.  It  is  best 
distinguished  from  the  other  varieties  by  albumen,  by  negative 
tests,  and  by  boiling  for  some  time  and  then  applying  the  tests 
for  the  ortho-acid. 

H3AsO4  and  H2CrO4  might  be  classified  with  this  group,  but 
a  little  care  on  the  part  of  the  student  will  detect  them  among 
the  bases,  and  a  little  thought  will  tell  him  how  to  prove 
whether  they  are  present  as  acids  or  bases. 


ACIDS.  77 


GROUP   III. 


ACETIC  ACID. 

VALERIANIC  ACID. 

STEARIC 

OLEIC 

LACTIC 

OXALIC 

SUCCINIC 

MALIC 


TARTARIC  ACID. 

CITRIC 

CARBOLIC        " 

BENZOIC 

SALICYLIC       « 

GALLIC 

TANNIC 


REACTIONS   OF  ACETIC   ACID   (HC2H3O2). 

1.  In  the  free  state  acetic  acid  is  readily  recognized  by  its 
odor. 

2.  H2SO4,  added  to  an  acetate  and  the  mixture  warmed, 
gives  the  characteristic  odor. 

3.  H2SO4   and    C2H5OH    in   equal   volumes   added   to   an 
acetate  form  acetic  ether,  readily  recognized  by  its  odor. 

4.  Fe2Cl6,  with  a  neutral   acetate,  forms  a  deep  red  color, 
due  to  ferric  acetate  —  Fe2(C2H3O2)6. 

VALERIANIC   ACID    (HC5H9O2). 

The  odor  is  sufficient  evidence  of  the  presence  of  this  acid.  This  odor  is 
developed  by  moisture  and  heat,  and  in  the  case  of  the  salts  by  the  addition 
of  H2S04. 

STEARIC   ACID 


Stearic  acid  is  a  white,  fatty  solid,  melting  at  70.5°  C.,  giving  when  combined 
with  potassium  a  soft  soap,  and  with  sodium  a  hard  soap  ;  from  both  it  separates 
as  an  oily  liquid  on  the  addition  of  HCl,  becoming  solid  on  cooling.  The  lead 
salt,  lead  stearate,  Pb(C18H35O2)2  is  insoluble  in  ether. 

OLEIC   ACID   (HC18H33O2). 

Oleic  acid  is  an  oily  liquid  at  ordinary  temperatures,  but  becomes  solid  at 
4°  C.  and  remains  so  until  the  temperature  rises  to  14°  C.,  when  it  again  be- 
comes liquid.  Lead  Oleate  —  Pb(C18H33O2)2,  prepared  by  neutralizing  sodium 
oleate  with  acetic  acid  and  adding  lead  acetate,  is  almost  entirely  soluble  in 
ether.  This  is  an  important  distinction  from  stearic  acid. 

LACTIC   ACID  (HC3H5O3). 

Lactic  acid  is  a  colorless,  syrupy  liquid,  of  a  slight  unpleasant  odor  and  a  very 
sour  taste,  soluble  in  water,  alcohol  and  ether,  but  insoluble  in  chloroform. 
When  heated  with  H2SO4,  CO  is  evolved.  Lactates  are  all  soluble  in  water, 
most  of  them  sparingly  ;  they  are  insoluble  in  ether. 

Hg2(NO3)2  boiled  with  strong  solution  of  a  lactate,  deposits  crimson  mer- 
curous  lactate  —  Hg2(C3H5O3j2. 


78  ANALYTICAL   CHEMISTRY. 

REACTIONS   OF   OXALIC   ACID  (H2C2O42H2O). 
Use  a  solution  of  ammonium  oxalate  ((NH4)2C2O4). 

1.  K2Mn2O8,  acidulated  with  H2SO4,  is  decolorized. 

2.  BaCl2  produces  a  white  precipitate  of  barium  oxalate — 
BaC2O4,  sparingly  soluble  in  acetic  or  oxalic  acid,  and  freely 
soluble  in  HC1,  HNO3  and  NH4CL 

3.  AgNO3   precipitates   white    silver    oxalate  —  Ag2C2O4, 
readily  soluble  in  hot  concentrated   HNO3,  sparingly  so  in 
dilute  acid,  soluble  in  NH4OH. 

4.  CaCl2  produces,  even  in  highly  diluted  solutions,  a  white 
precipitate  of  calcium  oxalate — CaC2O4,  soluble  in  HC1  and 
HNO3,  but  insoluble  in  acetic  acid. 

5.  FeSO4  with  dilute  solutions  causes  a  yellow  color;  with 
more  concentrated  solutions  and  warming,  a  yellow  precipitate 
ferrous  oxalate — FeC2O4,  falls.     This  precipitate  is  insoluble 
in  acetic  acid,  but  is  dissolved  by  HC1  and  HNO3. 

SUCCINIC   ACID  (H2C4H4O4). 
Use  a  solution  of  sodium  succinate  (Na2C4H4O4). 

1.  Bad 2  produces  no  precipitate,  but  on  the  addition  of  alcohol  a  white  pre- 
cipitate of  barium  succinate — BaC4H4O4,  falls,  soluble  in  NH4C1. 

2.  Fe2Cl6    causes   a  brownish-red,  bulky  precipitate  of  ferric  succinate — 
Fe2(C4H404)3. 

3.  Pb(C2H3O2)2  precipitates  white  amorphous  lead  succinate — PbC4H4O4, 
soluble  in  HNO3. 

MALIC   ACID  (H2C4H405). 

1.  CaCl2  fails  to  give  a  precipitate  with  the  acid  or  its  salts,  until  the  mixture 
is  boiled,  when  calcium  malate — CaC4H4O5H2O,  separates.     The  addition  of 
alcohol  will  also  cause  this  precipitate. 

2.  Pb(C2H3O2)2  precipitates  white  lead  malate— PbC4H4O53H2O. 

REACTIONS   OF   TARTARIC  ACID  (H2C4H4O6). 
Use  a  solution  of  sodium  tartrate  (Na2C4H4O6). 

1.  BaCl2   precipitates  white   barium   tartrate — BaC4H4O6, 
soluble  in  ammonium  salts,  and  in  HC1. 

2.  CaCl2  produces  a  white  precipitate  of  calcium  tartrate — 
CaC4H4O64H2O,  ammonium  salts  prevent  this  precipitation. 
On  adding   KOH  to  this  precipitate,  it  dissolves ;    boil  this 


ACIDS.  79 

solution,  calcium  tartrate  is  again  precipitated.     Acetic  acid 
is  also  a  solvent  for  this  precipitate. 

3.  AgNO3  causes  a  white  precipitate  of  silver  tartrate — 
Ag2C4H4O6;  on  adding  a  few  drops  of  NH4OH  to  dissolve  the 
precipitate,  and  boiling,  a  mirror  of  metallic  silver  forms  on 
the  test  tube.     The  test  tube  must  be  perfectly  cleaned  for  this 
reaction;  to  accomplish  this,  rinse  the  tube  with  KOH  solu- 
tion and  then  thoroughly  with  water. 

4.  Ca(OH)2  in  excess  causes  a  precipitate  of  calcium  tar- 
trate— CaC4H4O6.     This   precipitate  is    flocculent  at  first,  in 
which  state  it  is  readily  soluble  in  NH4C1,  but  after  standing  for 
some  time  it  becomes  crystalline,  in  which  case  it  is  insoluble 
in  NH4C1. 

5.  Tartaric  acid   and  tartrates  char  on   heating,  and  with 
H2SO4  the  odor  of  burnt  sugar  is  given  off. 

REACTIONS   OF   CITRIC  ACID  (H3C6H5O7H2O). 
Use  a  solution  of  sodium  citrate  (Na3C6H5O7). 

1.  BaCl2  produces  a  white  precipitate  of  barium  citrate — 
Ba3(C6H5O7)2,  soluble  in  excess  of  water,  in  ammonium  salts, 
and  in  acids. 

2.  CaCl2  precipitates  white  calcium  citrate — Ca3(C6H5O7)2, 
more  insoluble   in  hot  than   in   cold  water,  soluble  in  cold 
NH4C1,  but  insoluble  in  KOH. 

3.  AgNO3  produces  a  white  precipitate  of  silver  citrate  — 
Ag3C6H5O7;  on  boiling  no  metallic  mirror  is  formed. 

4.  Ca(OH)2,  in  excess,  does  not  cause  a   precipitate  until 
the  mixture  is  boiled.     This  is  an  important  distinction  from 
tartaric  acid. 

5.  Citric  acid  and  citrates  char  on  heating,  and  with  H2SO4 
give  off  the  odor  of  burnt  sugar. 

CARBOLIC  ACID  (C6H5OH). 

1.  HNO3,  with  an  aqueous  solution  of  the  acid,  forms  a  yellow  color,  due  to 
picric  acid— C6H2(NO2)3OH. 

2.  Fe2Cl6  produces  a  violet  blue  color. 

3.  A  piece  of  pine  wood  dipped  in  the  acid,  and  then  exposed  to  the  fumes  of 
HC1,  becomes  after  a  short  time  colored  blue. 

4.  Bromine  water  causes  a  white  flocculent  precipitate. 

5.  Albumen  is  coagulated  by  the  free  acid. 


80  ANALYTICAL   CHEMISTRY. 

BENZOIC  ACID  (HC7H5O2). 

npitates  from  neutral  solutions,  flesh-colored  ferric  benzoate — 
Fe2(C7H5O2)6,  soluble  in  HC1,  with  separation  of  benzoic  acid. 

2.  HC1   causes   the   separation  of  benzoic    acid   from    cold   solutions   of  the 
benzoates.  • 

3.  Bad 2  and  CaCl2  produce  no  precipitates  with  either  the  free  acid  or  its 
salts. 

SALICYLIC   ACID  (HC7H5O3). 

1.  Fe2Cl6  produces  a  deep  violet  color,  which  is  very  characteristic. 

2.  Warmed  with   HjSO^   and  methyl  alcohol,  the  fragrant  odor  of  methyl 
salicjfiate  (oil  of  winter  green)  is  developed. 

3.  HC1   causes  the    separation  of  the  free  acid  from   cold  solutions   of  the 
salicylates. 

GALLIC   ACID   (HC7H5O5.H2O). 

1.  FeSO4  produces  no  change. 

2.  Fe2Cl6  produces  a  bluish-black   precipitate,  which   dis- 
appears oh  heating. 

3.  KOH,  if  not  in  excess,  develops  slowly  a  deep  green 
color,  which  becomes  red  on  the  addition  of  acids.     Alka- 
line carbonates  cause  the  same  green  color,  although  more 
slowly. 

4.  No  precipitate  is  produced  with   either  gelatin   or  the 
alkaloids.     With  the  former,  however,  a  precipitation  takes 
place  when  gum  is  present. 

TANNIC   ACID   (C14H10O9). 

1.  FeSO4,  when  perfectly  pure,  causes  no  change;  in  the 
presence  of  oxygen,  however,  a  dark  color  rapidly  develops, 
which  on  standing,  slowly  becomes  a  precipitate. 

2.  Fe2Cl6  produces  a  bluish-black  precipitate. 

3.  Normal  solution  of  iodine  mixed  with  a  small  quantity 
of  amm«ia,  previously  diluted  with  ten  times  its  volume  of 
water,  jm)duces  a  brilliant  fed  color.     This  reaction  will  take 
placa  wJh  only  traces  of  tannin. 

4.  Gelatin  causes  a  white  flocculent  precipitate.     This  re- 
action  is   more  delicate  in  the 'presence   of  small   quantities 
of  aluni. 

5.  Alkaloids  produce  white  precipitates,  soluble  in  acetic 
acid  and  alcohol. 

6.  K(SbO)C4H4O6   causes   a   white,  gelatinous   precipitate. 
Most  metallic  salts  cause  precipitates  with  tannin. 


ACIDS. 


81 


DIRECTIONS  FOR  THE  DETECTION  OF  THE  ACIDS  IN  A  SOLU- 
TION CONTAINING  SOLUBLE  SALTS  OF  GROUP  III. 

As  many  of  the  acids  in  this  group  would  indicate  their 
presence  by  odor  or  physical  appearance,  a  method  of  sepa- 
rating only  the  more  important  and  closely  related  ones  will 
be  given.  The  list  will^therefore  be  limited  to  Acetic  acid,, 
Oxalic  acid,  Tartaric  acid,  Citric  acid. 

If  the  Solution  is  Acid  to  Litmus  Paper,  neutralize  with  NaOH. 

I.  To  a  small  portion  add  H2SO4  and  warm.     Acetic  acid 
will,  if  present,  be  detected  by  its  odor. 

II.  To  another  portion  add  NH4OH  until  slightly  alkaline, 
and  then  CaCl2 ;  allow  to  stand  (avoiding  heat)  for  ten  rruhutes, 
and  filter. 


Ppt.  H2C204,  H2C4H406. 
Wash,  pour  on  the  filter  HC2H3O2. 

Filt.  H3C6HB07. 
Boil  to  remove  NH4OH,  a  white  ppt. 
slowly  forms  on  sides  of  the  tube. 

• 

Ppt.  H2C204. 
Confirm  by  testing 
original  solution  with 
HC2H3O3andCaCl2. 

Filt.  H2C4H406. 
Add  NH4OH  until  slightly 
alkaline,  white  ppt. 
Confirm  by  forming  mirror  with 
AgNO3  in  original  solution. 

82  ANALYTICAL    CHEMISTRY. 

SECTION   III. 

DETECTION   OF   BASES   AND   ACIDS. 


SPECIAL   OBSERVATION. 

The  bases  must  always  be  determined  first  in  a  portion  of  the 
solution  or  powder,  according  to  chart,  page  65. 

If  only  the  alkali  metals  are  present  and  the  reaction  is 
neutral,  proceed  to  search  for  the  acids  according  to  DIVI- 
SION I. 

If  other  than  the  alkali  metals  are  present  and  the  reaction 
is  acid,  proceed  to  search  for  the  acids  according  to  DIVISION 
II. 

When  the  substance  is  not  entirely  dissolved  by  H2O,  HC1, 
or  a  mixture  of  HC1  and  HNO3,  proceed  for  both  bases  and 
acids  according  to  DIVISION  III. 

It  will  be  seen  from  the  above  that  before  beginning  the 
analysis  of  a  solution  it  is  necessary  to  try  its  action  on  litmus 
paper.  If  the  reaction  is  alkaline  it  is  necessary  to  neutralize 
with  HC1,  bearing  in  mind  the  precautions  on  page  66. 

When  the  analysis  of  a  solid  is  to  be  performed,  the  action 
on  litmus  paper  of  the  portion  soluble  in  water  is  also  to  be 
noted,  and  the  solution  made  neutral  if  necessary. 

All  the  physical  properties  of  the  solid  should  be  carefully 
observed. 

EXCEPTION  IN  SEARCHING  FOR  BASES. 
When  phosphates  are  present  in  acid  solution,  determined 
by  adding  a  few  drops  of  the  solution  to  some  HNO3  and 
(NH4)2MoO4  and  warming  gently,  the  following  chart  must  be 
used  in  the  separation  of  Group  IV.  That  is,  bring  the  pre- 
cipitate containing  Fe,  Ce,  Al,  Cr,  from  chart,  page  65,  and 
work  according  to  this.  Oxalates  also  cause  an  exception 
to  the  general  chart  for  separation  of  bases ;  when,  therefore, 
their  presence  is  suspected,  heat  the  substance  to  redness 
before  examining  for  bases. 


n  r 

?-T' 


DETECTION    OF    BASES    AND    ACIDS. 


83 


DIRECTIONS  FOR  THE  ANALYSIS  OF  GROUP   IV  WHEN   PHOSPHATES 
ARE   PRESENT. 

In  addition  to  Fe,  Ce,  Al,  Cr,  there  may  be  present  the  phosphates  of  Ca,  Sr,  Ba.  Mn,  Mg. 
Dissolve  the  precipitate  in  HC1,  add  Na2HP04  in  excess,  then  excess  of  NH4C2H3O2;  boil, 

filter. 


Ppt.  Fe2(P04)2  A12(P04)2 
Ce2(P04)2. 

Filt.  Ba,  Sr,  Ca,  Mn,  Mg,  Cr.    (If  green  Cr  is  present.) 
Add  K2CrO4)  warm,  filter. 

Wash  with  hot  H2O,  then  pour 

on  the  filter  hot  solution 

KOH. 

Ppt. 

Filt.  Sr,  Ca,  Mn,  Mg. 

Ba. 

Yellow. 

Add  very  dilute  H2SO4,  allow  to  stand,  filter. 
An  excess  of  H2SO4  should  be  avoided. 

Ppt.  Fe,  Ce. 
Wash,  dissolve  in 

Filt. 
A12(P04)2. 

HC1  and  test  a 
portion  with 
K4Fe(CNje 

Acidify  with 
HC2H3Oa 

white 

£• 

Confirm 

Filt.  Ca,  Mn,  Mg. 
Add  NH4C2H3O2  and  (NH4)2C2O4, 
filter. 

for  Fe. 

ppt.  =  Al. 

flby 

name 
test. 

Ppt. 
Ca. 
White. 

Filt.  Mg,  Mn. 
Evaporate  a  portion  to  dryness  and 
test  with  borax  bead  for  Mn. 

To  another  por- 
tion add 

H3C6H6O7, 
then  NH4OH  in 
excess  and 

To  the  remaining  portion  add 
Fe2Cle,  filter  out  the  Fe2(PO4)2, 
to  filtrate  add 

(NH4)2C204) 

NH4C1,  NH4OH  and  NH4HS, 

white 
precipitate  =  Ce. 

to  separate  Fa  and  Mn, 
filter  and  test  filtrate  with 

Na2HPO4,  white 

ppt.  =  Mg. 

DIVISION    I.—  When  the  bases  are  the  alkali  metals,  and  the  re- 
action is  neutral. 

RULE  I.  —  Evaporate  a  portion  of  the  solution  to  dryness,  and 
slowly  heat  to  redness. 

If  the  mass  chars,  one  or  more  of  the  following  organic 
acids  are  indicated  :  —  ACETATES,  TARTRATES,  CITRATES,  GAL- 
LATES,  TANNATES. 

NOTE.  —  Oxalates  do  not  char,  although  if  the  heating  take  place  slowly  a 
grayish  coloration  may  be  noticed,  the  residue  in  this  case  giving  off  CO2  on  the 
addition  of  H2SO4. 

Fe2Cl6  immediately  detects  TANNATES  and  GALLATES.  Gallic 
acid  crystallizes  in  the  cold  on  acidifying  the  solution.  Tan- 
nic  acid  precipitates  with  solution  of  gelatin.  The  other  acids 
are  separated  and  detected  according  to  chart,  page  81. 

RULE  II.  —  Add  to  a  second  portion  of  the  concentrated  solu- 
tion, or  the  dry  salt,  strong  H2SO4,  warm  gently,  and  note  any 
of  the  following  effects  :  — 

Effervescence  with  dilute  acid  in  the  cold,  \ 
no  odor  .........................................  J 


Carbonates. 


Effervescence  on  heating,  no  odor  ......  / 

\ 


Oxalates'    confirm 


CaC1      and 


HC2H302. 


84  ANALYTICAL   CHEMISTRY. 

Effervescence  with  dilute  acid  on  heating,  ) 

,        c  TJ  o  r  Sulphides, 

odor  of  H2S / 

Odor  of  SO2 Sulphites. 

Odor  of  SO2  with  precipitation  of  S Thiosulphates. 

c  Iodides,  confirm  by  starch  and  Cl 
Dark  brown  color  and  violet  fumes < 

i       ,  j.  ,    r  f  Bromides,   confirm    by   starch    and 

Dark  red  color  and  reddish  fumes 1 

I  Cl  water. 

Odor  of  HCN Cyanides. 

,.TT .-,,,     .,  ,,.       j         .  (   Ferro-    or    Ferri-cyanides,   confirm 

Odor  of  HCN  with  crystalline  deposit 1 

I  by  Fe2Cl6  and  FeSO4. 

Odor  of  acetic  acid Acetates. 

With  dilute  acid,  odor  of  chlorine Hypochlorites. 

With  strong  acid,  odor  of  chlorine Chlorates. 

Strongly  acid  suffocating  fumes Chlorides,  confirm  by  AgNO3. 

Strongly  acid  fumes  becoming  red  when  j    Nitrates,   confirm    by    FeSO4    and 
metallic  Cu  is  added \      H2SO4,  also  by  indigo  solution. 

c   Benzoates,  Succinates,  Valerianates, 

Characteristic  odors J 

(         Carbolates,  Hypophosphites. 

RULE  III. —  To  a  third  portion  add  BaCl2. 

A  white  precipitate  insoluble  in  HC1  indicates  SULPHATES. 

RULE  IV. — To  a  fourth  portion  add  CaCl2. 

A  white  precipitate  soluble  in  excess  of  H2O  indicates 
SULPHATES  ;  if  insoluble  in  excess  of  H2O  and  in  acetic  acid, 
OXALATES  are  indicated.  A  white  precipitate  soluble  in  KOH, 
reprecipitated  on  boiling,  indicates  TARTRATES;  confirm  by 
boiling  with  NH4OH  and  AgNO3,  forming  a  mirror  of  silver 
on  the  test  tube. 

A  white  precipitate  soluble  in  NH4C1  and  reprecipitated  on 
boiling  indicates  CITRATES.  Citrates  and  Tartrates  are  also 
detected  by  heating  some  of  the  dry  salt  with  H2SO4,  when 
the  odor  of  burnt  sugar  is  developed. 

When  all  four  of  these  acids  are  suspected  to  be  present, 
treat  the  solution  with  HC1  and  BaCl2  to  remove  sulphates ; 
then  neutralize  the  excess  of  HC1  with  NaOH,  and  proceed 
according  to  chart,  page  81. 

RULE  V.—  To  a  fifth  portion,  acidified  with  HNO3,  add  AgNO3. 

A  white,  curdy  precipitate,  immediately  and  completely 
soluble  in  NH4OH  indicates  CHLORIDES.  If  slowly  soluble 
in  NH4OH,  BROMIDES  may  be  present;  confirm  by  chlorine 
water  and  starch. 


DETECTION    OF    BASES    AND    ACIDS.  85 

A  yellowish  precipitate  insoluble  in  NH4OH  =  IODIDES. 
A  white  precipitate  soluble  in  strong,  hot  HNO3  =  CYA- 
NIDES. When  bromides  and  iodides  are  present  they  may 
be  separated,  by  adding  to  the  original  solution  chlorine 
water  and  starch ;  continue  the  addition  of  chlorine  water 
until  the  blue  color  of  the  starch  iodide  is  discharged,  then 
shake  with  chloroform  ;  when  a  red  color  is  imparted  to  the 
separated  chloroform,  Br  is  indicated. 

For  further  instruction  regarding  the  separation  of  these 
acids,  consult  chart,  page  69. 

RULE  VI. — Add  to  a  sixth  portion  of  the  solution  magnesia 
mixture. 

A  white  precipitate  indicates  PHOSPHATES  and  ARSENATES. 
In  another  portion  separate  the  As  by  HC1  and  H2S,  and 
repeat  the  reaction  for  H3PO4  in  the  filtrate  by  neutralizing 
and  adding  magnesia  mixture. 

RULE  VII. —  To  a  seventh  portion  add  Fe2Cl6. 
This  reagent  readily  indicates  TANNATES,  GALLATES,  FERRO- 
and  FERRI-CYANIDES. 

RULE  VIII. —  To  an  eighth  portion  apply  the  special  test  with 
H2SO4  and  FeSO4/<?r  NITRATES.  And  apply  the  special  test 
for  BORATES,  by  adding  H2SO4,  alcohol,  and  igniting  to  obtain 
the  green  flame. 

DIVISION  II. — When  the  solution  is  acid,  or,  if  a  solid,  requires  the 
use  of  HC1  to  dissolve  it. 

After  determining  the  bases,  bearing  in  mind  the  exception  in 
regard  to  phosphates,  a  portion  of  the  substance  is  tested  accord- 
ing to  Rules  I  and  II,  DIVISION  I. 

A  second  portion  is  treated  as  follows  : — 

Boil  with  strong  solution  Na2CO3  ;  filter. 


Ppt.   consists  of  the 
carbonates,  hydrates, 
and   oxides,   of  the 
bases  present,  and 
is  disregarded. 

Filt.  contains  the  acids  in  combination  with  Na. 
Divide  in  two  parts. 

First  portion. 
Add  HC1  until  slightly  acid, 
boil  to  remove  CO2,  neu- 
tralize with  NaOH,  and 
test  according  to  Rules 
III,  IV,  VI,  VII,  and  VIII, 
DIVISION  I. 

Second  portion. 
Add  HNO3  until  slightly  acid, 
boil  to  remove  CO2,  neutral- 
ize with    NaOH,   and  test 
according  to    Rule   V, 
and    if   necessary, 
III,  IV,  VI,  VII, 
DIVISION  I. 

86  ANALYTICAL   CHEMISTRY. 

DIVISION    III. — When   the    substance   is    not    entirely   dissolved   by 
H2O,  HC1,  or  a  mixture  of  HC1  and  HNO3. 

Add  water  and  filter ;  the  filtrate,  if  it  contains  solid  matter, 
is  examined  for  bases  by  chart,  page  65,  and  for  acids  by 
Division  I.  The  residue  after  treatment  with  water  is  treated 
with  HC1,  or  if  necessary,  with  a  mixture  of  HC1  and  HNO3, 
and,  if  it  contain  any  solid  matter,  examined  for  bases  by 
chart,  page  65,  and  for  acids  according  to  Division  II.  The 
insoluble  residue  consists  of  one  or  more  of  the  following 
substances :  Sulphates  of  Ba,  Sr,  Ca,  Pb ;  Chlorides,  Bromides 
and  Iodides  of  Ag,  Pb  ;  certain  Oxides  which  have  been  highly 
heated,  as  Fe2O3,  A12O3,  Cr2O3  and  SnO2 ;  Silica  and  Silicates ; 
Carbon  and  Sulphur  (the  last  two  are  readily  detected  in  the 
preliminary  examination  by  heat).  This  insoluble  residue  is 
mixed  with  about  four  times  its  weight  of  dry  sodium  car- 
bonate and  fused  on  a  platinum  foil ;  the  mass  is  boiled  with 
water,  and  the  soluble  portion  examined  for  bases  by  chart, 
page  65,  and  for  acids  according  to  Division  I.  The  in- 
soluble portion  is  treated  with  HC1,  evaporated  to  dryness, 
redissolved  in  water  with  a  little  HC1,  and  examined,  after 
determining  the  bases,  according  to  Division  II.  Any  insolu- 
ble substance  remaining  is  probably  silica,  which  may  be 
determined  by  a  bead  of  salt  of  phosphorus  before  the  blow- 
pipe, showing  the  skeleton  of  silica.  If  the  metals  lead  and 
silver  are  present  in  this  insoluble  residue,  they  are  best  deter- 
mined by  treating  the  residue  insoluble  in  HC1,  by  HNO3  and 
examining  according  to  Division  II.  Many  of  these  insoluble 
substances  are  detected  in  a  preliminary  examination  by  the 
use  of  the  blowpipe,  as  explained  in  the  tests  given  under  the 
individual  elements. 

The  following  chart,  taken  largely  from  Dr.  Miiter's  Ana- 
lytical Chemistry,  will  be  found  useful  to  consult  after  deter- 
mining the  bases  and  before  commencing  the  examination  for 
acids. 


DETECTION   OF   BASES   AND   ACIDS. 


87 


Bases 
found. 

If  soluble  in  water  test  for  the 
following  acids. 

If  insoluble  in  water,  but 
soluble  in  acids,  test  for 
the  following  acids. 

If  insoluble 
in  acids,  fuse 
with  Na2CO3 
and  test  for, 

Ag, 

HNO3,  HNO2,  H2SO4,  HC2H3O2, 

Oxide,  Sulphide,  H2CO3, 
H3PO4,  HCN,  H2C2O4, 
H2C4H406,  H3C6H5Cy 

Cl,  I,  Br. 

Hg(ous), 

HNO3,  HC2H3O2,  H2SO4, 

Oxide,  Sulphide,  Cl,  1, 
Oxysulphate. 

Cl,  I,  Br. 

Hg(ic), 

Cl,  HNO3,  H2SO4,  HC2H3O2, 

Oxide,  Sulphide,  I, 
Oxysulphate. 

Sulphide, 
Iodide. 

Pb, 

HC2H3O2,  HNO3, 

Oxide,  Sulphide,  H2CO3, 
H3PO4,  H2C2O4. 

H2SO4,  Cl, 
I,  H2CrO4. 

Bi, 

HNO3,  Cl,  H2SO4,  HC2H3O2, 

Oxynitrate,  Oxychloride, 
Oxide,  Sulphide,  H2CO3, 
H3P04. 

None. 

Cu(ic), 

Cl,  HNO3,  H2SO4,  HC2H3O2, 

Oxide,  Sulphide,  H2CO3, 
H3PO4,  Oxyacetate. 

" 

Cu(ous), 

H2S04, 

Oxide,  I. 

" 

Cd, 

Cl,  HN03,  I,  H2S04, 

Oxide,  Sulphide,  H2CO3 
H3P04. 

u 

Sb, 

Cl,  H2C4H406. 

Oxide,  Sulphide, 
Oxychloride. 

K 

Sn(ic), 

Cl, 

Oxide. 

" 

Sn(ous), 

Cl,  H2SO4, 

Oxide,  Sulphide,  H3PO4 
H2CrO4. 

It 

Au, 

Cl, 

Sulphide. 

" 

Pt, 

Cl, 

Sulphide. 

11 

Fe(ic), 

Cl,  HNO3,  H2SO4,  HC2H3O2, 

Oxide,  Sulphide,  I, 
H3P04. 

« 

Fe(ous), 

Cl,  H2S04,  I, 

Oxide,  Sulphide,  H2CO3 
H3P04. 

« 

Al, 

Like  Iron, 

Oxide,  H3PO4. 

«« 

Ce, 

Cl, 

Oxide,  H2C2O4. 

Oxide. 

Cr, 

Cl,  H2S()4,  HC,H302,  HN03, 

Oxide,  H3PO4. 

Oxide. 

Mn, 

Like  Chromium, 

Oxide,  Sulphide,  H2CO3 
H3P04. 

None. 

Zn, 

Like  Chromium, 

Like  Manganese. 

" 

Co, 

Cl,  HN03,  H2S04, 

Oxide,  Sulphide,  H2CO3, 
H3P04. 

u 

Ni, 

Like  Cobalt, 

Like  Cobalt. 

" 

Ba, 

Cl,  HNO3,  HC2H3O2, 

H2C03,  H3P04,  H2C204 
H2CrO4. 

Sulphate. 

Sr, 

Like  Barium, 

Like  Barium. 

« 

Ca, 

Like  Barium, 

Like  Barium. 

Sulphate. 

Mg, 

Oxide,  Cl,  H2SO4, 

H2CO3,  H3PO4. 

None. 

Li, 

All  Radicals, 

H3P04. 

«< 

K, 

All  Radicals, 

K2PtCl6,  KHC4H406. 

" 

Na, 

All  Radicals, 

None. 

" 

NH4, 

All  Radicals, 

None. 

a 

88  ANALYTICAL   CHEMISTRY^ 


SECTION  IV. 

SOME    OF    THE    REACTIONS    AND     TESTS    OF    PURITY 
OF  THE   MORE  IMPORTANT  ORGANIC   COMPOUNDS. 


CHLOROFORM  (CHC1,). 

"  A  heavy,  clear,  colorless,  diffusive  liquid,  of  a  character- 
istic, pleasant,  ethereal  odor,  a  burning,  sweet  taste,  and  a 
neutral  reaction." 

1.  Agitate  about  5  c.  c.  with  twice  its  volume  of  distilled 
water,    separate   and   test   the   aqueous   portion  with    litmus 
paper ;  no  change  should  be  produced,  showing  the  absence 
of  acids.     To  a  small  portion  of  the  water  add  AgNO3;   a 
white  precipitate  would  indicate   HC1,  the  result  of  decom- 
position.    To  another  portion  add  solution  of  KI  ;  a  reddish 
color  would  indicate  Chlorine.     Another  portion  warmed  with 
solution  KOH  should  not  become  colored,  showing  the  ab- 
sence of  aldehyde. 

2.  10  c.  c.  of  chloroform  mixed  with  half  its   volume   of 
strong  H2SO4,  and,  after  agitation,  set  aside,  should  not  cause 
a  color  in  either  liquid  for  twenty-four  hours. 

3.  H2SO4   and    K2Cr2O7,  mixed  with    an   equal  volume   of 
chloroform  and  warmed,  will  become  green,  indicating  the 
presence  of  a  small  quantity  of  alcohol. 

IODOFORM  (CHIj). 

lodoform  occurs  in  small,  lemon-yellow,  lustrous  crystals, 
having  a  saffron-like  odor,  and  an  unpleasant,  slightly  sweetish, 
iodine-like  taste.  Its  solutions  have  a  neutral  reaction. 

1.  Digest  a  small  quantity  of  iodoform  with  alcoholic  solu- 
tion of  KOH,  acidify  with  dilute  HNO8,  and  add  starch  paste, 
when  the  blue  color  of  starch  iodide  will  be  developed. 

2.  Agitate  a  small  quantity  of  iodoform  with  distilled  water 
and  filter;  the  filtrate  should  not  affect  litmus  paper.     Add  to 
a  portion  of  the  filtrate  AgNO3 ;  a  yellowish  precipitate  would 
indicate   iodides  as  an  impurity.      Evaporate   another  small 


REACTIONS   AND   TESTS   OF   ORGANIC    COMPOUNDS.  89 

portion  to  dryness ;  no  residue  should   remain,  indicating  the 
absence  of  soluble  impurities. 


ALCOHOL  (C2H5OH). 

"  A  transparent,  colorless,  mobile  and  volatile  liquid,  of  a 
characteristic,  pungent    and    agreeable  odor,  and  a   burning 
taste.    It  should  not  change  the  color  of  blue  litmus  paper,  pre 
viously  moistened  with  water.    It  boils  at  78°  C.  and  is  readily 
inflammable,  giving  a  blue  flame  without  smoke." 

1.  Add  to  5  c.  c.  of  alcohol  in  a  test  tube  an  equal  volume 
of  H2SO4  to  which  has  been  added  a  small  quantity  of  K2Cr2O7, 
and  warm  ;  the  odor  of  aldehyde  will  be  developed,  the  liquid 
in  the  tube  becoming  green. 

2.  On  heating   10  c.  c.  of  alcohol  with  about  one-fourth  its 
volume  of  H2SO4  the  odor  of  ether  will  be  evolved. 

3.  Equal  volumes  of  H2SO4,  acetic  acid  and   alcohol,  on 
warming,  will  give  the  odor  of  acetic  ether.     If  much  amyl 
alcohol  be  present  the   odor   of   amyl  acetate   may  also  be 
recognized. 

4.  Evaporate  10  c.  c.  of  alcohol  to  about  one-fifth  its  bulk, 
and  add  an-  equal  volume  of  strong  sulphuric  acid  ;  a  reddish 
color  will  indicate  amyl  alcohol. 

5.  Mix  in  a  test  tube  5  c.  c.  of  alcohol  with  an  equal  volume 
of  solution  KOH ;  an  immediate  darkening  will  indicate  the 
presence  of  methyl  alcohol,  aldehyde  or  oak  tannin. 

CHLORAL  HYDRATE  (C2C13OH,H2O). 

(Chloral,  U.  S.  P.) 

Chloral  hydrate  occurs  in  colorless,  transparent  crystals, 
slowly  vaporizing  on  exposure  to  the  air,  having  an  aromatic, 
penetrating  and  slightly  acrid  odor,  a  bitterish  caustic  taste 
and  a  neutral  reaction.  It  liquefies  when  mixed  with  carbolic 
acid  or  with  camphor. 

1.  In  a  test  tube  heat  to  boiling  a  small  quantity  of  chloral 
hydrate  and  water  and  add  solution  KOH  ;  a  vaporous  milky 
mixture  of  chloroform   results,  readily  recognized  by  its  odor, 
while  formate  of  potassium  remains  in  solution. 

2.  To  a  hot  aqueous  solution  of  chloral   hydrate  in  a  test 


90  ANALYTICAL   CHEMISTRY. 

tube,  add  AgNO3,  then  NH4OH  and  boil ;  a  silver  mirror  is 
deposited  on  the  tube. 

3.  To  another  portion  of  an  aqueous  solution  add  a  few 
drops  of  HNO3  and  AgNO3 ;   no  precipitate  should  be  pro- 
duced, indicating  the  absence  of  HC1. 

4.  Add  NH4HS  to  an  aqueous  solution ;  a  reddish-brown 
coloration  results,  which  on  standing  deposits  a  reddish-brown 
compound  mixed  with  sulphur. 

5.  A  crystal  of  chloral   hydrate   heated  on   platinum   foil 
should  volatilize  without  residue. 

GLYCERIN  (C3H5(OH)3). 

"  A  clear,  colorless  liquid,  of  a  syrupy  consistence,  oily  to 
the  touch,  hygroscopic,  without  odor,  very  sweet  and  slightly 
warm  to  the  taste,  and  neutral  in  reaction." 

1.  Make  a  borax   bead   in   the   loop    of  a   platinum  wire, 
moisten  it  with  a  small  quantity  of  glycerin  (previously  ren- 
dered alkaline  with  a  dilute  solution  of  soda),  and  hold  in  the 
colorless  flame  of  a  Bunsen  burner  ;  the  flame  is  tinged  green, 
owing  to  the  liberation  of  boric  oxide. 

2.  Warm  5  c.c.  of  glycerin  with  an  equal  volume  of  H2SO4;  no 
coloration  should  result,  showing  the  absence  of  cane  sugar. 

3.  Heat  a  few  drops  on  platinum   foil ;  no  residue   should 
be  left. 

4.  A  portion  heated  nearly  to   the  boiling  point  with  an 
equal  volume  of  Fehling's  solution,  should  not  deposit  a  red 
precipitate  of  cuprous  oxide,  Cu2O.     The  same  result  should 
be  obtained  if  the  glycerin  be  previously  boiled  with  a  small 
quantity  of  HC1,  showing  the  absence   of  sugars,  starch   and 
dextrin. 

GLUCOSE  (C6H1206). 
(Dextrose,  Grape  Sugar?) 

1.  To  a  dilute  solution  of  glucose,  add  an  equal  volume  of 
Fehling's  solution  and  boil ;  a  red  precipitate  of  cuprous  oxide, 
Cu2O,  is  deposited. 

2.  A  strong  solution  mixed  with  H2SO4  in  the  cold ;  no 
change   should    occur;    if,  however,  organic    impurities    are 
present,  a  dark  coloration  will  result. 


REACTIONS   AND   TESTS   OF   ORGANIC    COMPOUNDS.  91 

3.  To  a  strong  solution  of  glucose  add  a  solution  of  KOH 
and  warm ;  the  solution  becomes  yellow  and  finally  brown. 

SACCHAROSE  (C^H^Ou). 
(Cane  Sugar?) 

1.  Boil  a  moderately  dilute   solution  of  sugar  with  a  few 
drops    of   HC1,   neutralize  and  add   Fehling's    solution ;  red 
cuprous   oxide,  Cu2O,  will  be  precipitated.     If  the  sugar  be 
pure  this  reaction  will  not  take  place  unless   it  be  previously 
boiled  with  the  acid. 

2.  To  a  cold   saturated  solution  add  an  equal  volume   of 
concentrated  H2SO4;  there   is  an   immediate  blackening  and 
swelling  until  the  tube  is  filled  with  a  dry  coke. 

AMYLOSE  (C6H1005)X. 

(Starch.) 

Starch  usually  occurs  as  a  fine,  white  powder,  insoluble  in 
.cold  water,  soluble  in  boiling  water,  forming  a  thick  paste 
when  cold. 

1.  To  a  cold  dilute  solution  add  potassium  iodide  and  a  few 
drops  of  chlorine  water;  the  deep  blue  color  of  starch  iodide 
results.     This   is  decolorized   by   heating   and   reappears    on 
cooling.     It  is  bleached  by  excess  of  chlorine  water. 

2.  Heat  a   small  quantity  of  starch  on  a  platinum   foil ;  it 
should  char  and  finally  disappear  without  residue. 

3.  Triturate  a  small   quantity  of   starch   with    cold    water, 
filter,  and   test  the   filtrate  with    litmus ;  no    change    should 
take  place. 

MORPHINE  (C17H19NO3,H2O). 

1.  To  some  crystals  of  morphine  sulphate  on  a  crucible  lid, 
or  other  white   porcelain    surface,    add   a    small    quantity  of 
neutral  solution  of  ferric  chloride ;  a  blue  color  is  produced, 
which    rapidly  fades.     The   presence  of  free    sulphuric   acid 
must  be  avoided. 

2.  Similarly,  to  another  portion  add  a  drop  of  strong  HNO3; 
a  red  color  at  first  forms,  which  rapidly  becomes  yellow. 

3.  To  a  solution  of  the   sulphate   add   carefully  a   dilute 


92  ANALYTICAL   CHEMISTRY. 

solution  of  KOH  ;  a  white  precipitate  forms,  which  is  readily 
soluble  in  excess.  NH4OH  and  Na2CO3  produce  the  same 
precipitate,  insoluble  in  excess. 

4.  A  few  crystals  on  a  white  porcelain  surface  should 
entirely  dissolve  in  strong  H2SO4  without  color,  becoming 
reddish  on  standing,  and  on  the  addition  of  a  crystal  of 
K2Cr2O7  a  greenish  color  is  produced.  A  purple  or  violet 
color  would  indicate  the  presence  of  strychnine  or  brucine. 

STRYCHNINE  (C21H22N2O2). 

1.  A  small  quantity  of  the  alkaloid  dissolved  in  a  drop  or 
two  of  H2SO4,  and  a  small  crystal  of  K2Cr2O7  added  will  cause 
a  deep  blue  color,  rapidly  changing  to  violet,  then  cherry-red, 
and  finally  fading. 

2.  Strong   HNO3  added   to  strychnine   should   not   cause 
more   than   a  very  faint  red   color,  showing   the   absence  of 
brucine. 

3.  KOH  added  to  a  solution  of  the  sulphate  causes  a  white 
precipitate,  insoluble  in  excess. 

QUININE  ((^HMN2Oa,3H20). 

1.  To  an  aqueous  solution  of  the  sulphate  add  fresh  chlorine 
water  and  then  NH4OH  in  slight  excess,  when  a  green  color  is 
formed,  due  to  Thalleioquin. 

2.  To   a  very  dilute,  slightly  acid   solution,   add   chlorine 
water,  then  a  small  quantity  of  K4Fe(CN)6,  and  finally  a  few 
drops  of   NH4OH  ;    a  red   color  is  produced,  which  rapidly 
fades. 

3.  Add   KOH  to  an  aqueous  solution  of   the   sulphate ;    a 
white  precipitate  forms,  insoluble  in  excess.     Under  the  same 
circumstances  NH4OH  produces  a  white  precipitate,  soluble  in 
excess. 

CINCHONIDINE  (C^H^O). 

1.  To  an  aqueous  solution  of  the  sulphate  add  NH4OH  ;  a 
white  precipitate  forms,  almost  insoluble  in  excess. 

2.  KNaC4H4O6  added  to  a  neutral  solution  produces  a  white 
precipitate. 


REACTIONS   AND   TESTS    OF    ORGANIC    COMPOUNDS.  93 

3.  To  a  cold  saturated  solution  of  the  sulphate  add  KNa- 
C4H4O6  in  slight  excess,  allow  to  stand  a  short  time  at  about 
15°  C,  filter  and  to  the  filtrate  add  a  drop  of  NH4OH ;  not 
more  than  a  slight  turbidity  should  result,  showing  the  absence 
of  more  than  small  quantities  of  the  other  cinchona  alkaloids. 

CINCHONINE  (C20H24N2O). 

1.  To  a   solution  of  the    sulphate   add    NH4OH ;  a  white 
precipitate  results,  insoluble  in  excess. 

2.  To  a  saturated  solution   add    chlorine  water   and  then 
NH4OH  in  slight  excess;  a  white  precipitate  is  formed,  free 
from  green  color. 

ATROPINE  (C17H22NO3). 

1.  Add  to  a  small  quantity  of  the  alkaloid  a  drop  of  H2SO4, 
no  change  occurs  ;  add  to  the  mixture  HNO3,  no  color  is  pro- 
duced, showing  the  absence  of  and  difference  from  morphine  ; 
similarly  when  a  crystal  of  K2Cr2O7  is  added  no  change  occurs, 
indicating  the  absence  of  and  difference  from  strychnine. 

2.  To  a  small  quantity  of  atropine  in  a  test  tube  add  H2SO4 
and  heat  until  it  turns  brown  ;  dilute  with  water  and  boil,  when 
the  characteristic  odor  of  orange  flowers  will  be  developed. 

3.  Boil  with  K2Cr2O7  and  dilute   H2SO4,  and  add  KOH  in 
excess,  when  a  herring-like  odor  will  be  given  off. 

CAFFEINE  (C8H10N402,H20). 

1.  To  an  aqueous  solution  of  the  alkaloid  add  a  solution  of 
potassio-mercuric  iodide  ;  no  precipitation  occurs,  indicating 
the  absence  of  and  distinction  from  most  other  alkaloids. 

2.  Treat  a  small  quantity  with  chlorine  water  and  evaporate 
to  dryness,  when  a  yellow  mass  will  remain,  which  on  moisten- 
ing with  NH4OH  will  become  purple. 

VERATRINE. 

I.  To  a  small  quantity  on  a  crucible  lid,  add  a  drop  of 
HNO3;  it  dissolves  with  a  yellow  color,  which  soon  passes  to 
a  reddish-yellow,  then  to  an  intense  scarlet  and  finally  to  a 
violet-red. 


94  ANALYTICAL   CHEMISTRY. 

2.  Similarly,  to  another  portion  add  H2SO4;  very  little  change 
takes  place,  but  on  warming  a  resinous  mass  is  formed,  which 
dissolves  with  a  deep  red  color. 

3.  Repeat  the  above  reaction,  using  HC1,  when  a  blood-red 
color  results. 

SALICIN  (C13H1807). 

1.  Add  cold,  concentrated  H2SO4  when  a  red  color  is  pro- 
duced, which  disappears  on  adding  water,  a  dark  red  powder 
insoluble  in  water  and  in  alcohol  being  deposited. 

2.  The  aqueous  solution  of  salicin   should   not  be   precipi- 
tated by  tannic  or  picric  acids,  nor  potassio-mercuric  iodide, 
indicating  the  absence  of  and  difference  from  most  alkaloids. 

SANTONIN  (C15H1803). 

1.  To  an  alcoholic  solution  of  KOH  add  santonin;  it  will 
dissolve  with  a  red  color,  gradually  becoming  colorless. 

2.  To  another  portion  of  santonin  add  cold  concentrated 
H2SO4 ;  there  is  gradually  formed  a  yellow  color,  becoming 
red  and  finally  brown. 

3.  Santonin  should  give  no  precipitates  with  the  ordinary 
reagents  for  the  alkaloids. 

NOTE. — All  the  preceding  organic  compounds  may  be  tested 
for  mineral  impurities  by  heating  on  a  platinum  foil ;  no  residue 
should  be  left. 


PART  THIRD. 


QUANTITATIVE  ANALYSIS 


PART  III.— QUANTITATIVE    CHEMICAL 
ANALYSIS. 

SECTI0N  I. 

GRAVIMETRIC    ESTIMATION. 

PRELIMINARY  DIRECTIONS. 

The  following  course  supposes  the  student  to  be  familiar 
with  the  operations  in  qualitative  analysis,  as  well  as  with  the 
process  of  weighing,  or  in  a  position  to  be  instructed  by  some 
one  familiar  with  the  use  of  analytical  balances  and  weights. 
The  following  rules,  however,  should  be  kept  constantly  before 
those  using  a  fine  balance  : — 

1 .  Never  put  any  chemical  substance  directly  on  the  pan  of  a 
balance,  but  always  in  a  clean,  dry  watch  crystal. 

2.  Never  put  on  or  take  off  a  weight,  or  anything  else,  from  a 
balance  when  it  is  resting  on  the  knife  edges. 

3.  All  volatile  acids  and  other  corrosive  substances  should  be 
weighed  in  stoppered  weighing  tubes. 

4.  Do  not  keep  the  balance  open  longer  than  is  absolutely 
neccessary. 

PRECIPITATION. 

This  should  be  conducted  in  beaker  glasses.  Avoid  a  large 
excess  of  the  reagent,  except  in  a  few  cases  where  it  is 
directed.  The  filtrate  should  always  be  tested  with  a  few 
drops  more  of  the  reagent,  to  determine  that  the  precipitation 
has  been  complete.  Distilled  water  is  used  in  all  quantitative 
determinations. 

FILTRATION. 

A  filter  paper  having  a  very  small  ash  should  be  used ;  the 
Swedish  is  probably  the  best.  Plain  filters  are  always  used, 
made  by  folding  a  circular  piece  of  the  paper  twice,  and  open- 
ing so  as  to  form  a  cone,  which  will  fit  exactly  into  a  funnel, 
and  allow  all  the  liquid  to  pass  through  the  precipitate  and 
7  97 


98  QUANTITATIVE   ANALYSIS. 

escape  at  the  lower  point  into  the  neck  of  the  funnel.  The 
precipitate  should  be  allowed  to  settle  in  the  beaker  after 
precipitation ;  the  clear  supernatant  liquid  is  then  poured  on 
the  filter,  using  a  glass  rod  against  the  edge  of  the  beaker,  to 
conduct  it  into  the  funnel  without  loss.  The  precipitate  is 
then  washed  into  the  filter  with  the  aid  of  the  wash  bottle, 
and  the  washing  continued,  until  by  the  appropriate  test  it  is 
found  to  be  complete.  The  washing,  in  many  cases,  is  to  be 
done  with  hot  water. 

DRYING. 

The  filter  and  its  contents  are  usually  dried  in  the  funnel 
when  intended  for  ignition,  but  if  to  be  weighed  without  burn- 
ing, it  is  transferred  to  two  watch  crystals  clamped  together, 
and  dried  in  an  air  bath,  at  from  120°  to  130°  C,  until  it  ceases 
to  lose  weight. 

IGNITION. 

This  is  usually  done  in  a  platinum  or  porcelain  crucible,  the 
precipitate  being  separated  as  much  as  possible  from  the  filter, 
and  the  latter  burned  on  the  lid,  then  placed  in  the  crucible 
with  the  precipitate  and  the  whole  brought  to  a  low  or  bright 
red  heat,  as  the  case  requires.  It  is  .then  cooled  in  a  desiccator, 
until  nearly  the  temperature  of  the  balance  room,  when  it  is 
ready  for  weighing. 


Examples  for  Practice   by  Gravimetric  Estimation. 

BARIUM  CHLORIDE  (BaCl22H2O). 

Estimation  of  the  Barium. 

Weigh  about  .500  gram  of  the  pure  salt,  dissolve  in  100  c.c. 
of  water  in  a  beaker  of  about  200  c.  c.  capacity.  Add  dilute 
sulphuric  acid,  drop  by  drop,  as  long  as  any  precipitate  is 
produced,  and  then  boil  the  mixture.  Allow  the  precipitate  to 
settle,  pour  the  clear  supernatant  liquid  into  a  filter,  boil  with 
a  fresh  quantity  of  water,  and  after  pouring  in  the  clear  liquid, 
transfer  the  precipitate  to  the  filter.  Remove  the  small  part 
•of  the  precipitate  which  adheres  to  the  sides  of  the  beaker 
with  a  glass  rod  having  a  short  section  of  rubber  tubing 
slipped  over  the  end.  Wash  the  precipitate  with  hot  water, 


GRAVIMETRIC    ESTIMATION.  99 

until  the  washings  cease  to  cause  a  turbidity  with  solution  of 
barium  chloride,  then  dry,  ignite  and  weigh  as  barium  sul- 
phate, BaSO4. 

The  following  calculation  will  serve  to  illustrate  the  general 
method : — 

Ba  =  136.8  =  56.16  per  cent. 
C12  =    70.8  =  29.06       " 
2H2O  =    36.    =  14.78       " 

243.6      100.00 

Supposing  we  take  .365  gram  of  barium  chloride  and  find 
•349  gram  of  barium  sulphate,  we  have  as — 

Molecular  wt.  Atomic  wt.  Wt.  of  BaSO4  Wt.  of  Ba 

of  BaSO4.  of  Ba.  found.  found. 

232.8  :  136.8  :     :  .349  :  .205  -f 

Wt.  of  BaCl22H2O  Wt.  of  Ba  per  cent,  of  Ba 

taken.  found.  found. 

As  .365  :  .205  -f  :     :  100  :  56.16. 


ESTIMATION  OF  THE  CHLORINE. 
To  accomplish  this  weigh  another  portion  of  the  salt, 
dissolve  in  about  100  c.  c.  of  water,  acidify  with  HNO3,  add 
AgNO3  in  slight  excess  and  heat  to  the  boiling  point.  Stir 
the  mixture  while  hot,  until  the  precipitate  coagulates,  allow 
to  settle,  decant  the  clear  liquid  to  a  filter,  then  rapidly  transfer 
the  precipitate,  and  wash  it  well  with  hot  water,  dry  and  ignite, 
burning  the  filter  separately.  If  an  appreciable  amount  of  the 
precipitate  adheres  to  the  filter  it  is  best  to  moisten  the  residue 
after  burning  with  a  drop  or  two  of  nitrohydrochloric  acid,  to 
reconvert  any  metallic  silver  which  may  have  formed  into 
chloride.  The  whole  operation  should  be  conducted  as  rapidly 
as  possible,  as  silver  chloride  is  acted  on  by  the  light.  The 
calculation  is  performed  precisely  like  that  for  barium,  substi- 
tuting C12  for  Ba,  and  AgCl  for  BaSO4. 

COPPER  SULPHATE  (CuSO4sH2O). 

Estimation  of  the  Copper. 

Weigh  from  .500  to  I  gram,  dissolve  in  a  small  quantity 
of  water  in  a  porcelain  dish ;  add  potassium  hydrate  in  slight 
excess  and  boil,  to  convert  the  cupric  hydrate  first  formed  into 
cupric  oxide.  Wash  two  or  three  times  by  decantation,  then 


100  QUANTITATIVE   ANALYSIS. 

transfer  to  the  filter  and  again  wash  with  hot  water  until  the 
filtrate  ceases  to  change  red  litmus  paper.  The  precipitate  is 
separated  from  the  filter,  and  the  latter  after  burning  is  moist- 
ened with  a  drop  of  nitric  acid  and  again  heated,  to  reconvert 
any  metallic  copper  which  may  have  been  formed ;  this  residue 
is  then  transferred  to  the  crucible  with  the  precipitate  and  the 
whole  ignited  and  weighed  as  cupric  oxide  CuO. 

ESTIMATION  OF  SULPHURIC  ACID. 
Weigh  another  portion  of  the  copper  sulphate,  dissolve  and 
precipitate  the  sulphuric  acid  with  barium  chloride.     Treat  the 
precipitate  exactly  as  in  the  estimation  of  barium. 

POTASSIUM  NITRATE  (KNO3). 

Estimation  of  the  Potassium. 

To  a  weighed  portion  of  the  salt  dissolved  in  a  small 
quantity  of  water,  add  hydrochloric  acid,  and  evaporate  to 
dryness,  dissolve  in  a  small  quantity  of  water,  add  platinic 
chloride  in  excess  ana  evaporate  nearly  to  dryness  on  a  water 
bath,  keeping  the  water  in  the  bath  just  below  the  boiling 
point.  Add  80  per  cent,  alcohol  to  the  residue,  allow  to  stand 
some  time,  then  transfer  to  a  small  weighed  filter,  and  wash 
repeatedly  with  small  quantities  of  the  same  strength  alcohol. 
Dry  the  precipitate  first  in  the  air,  then  in  the  air  bath  at  130° 
C.,  and  weigh  between  two  watch  crystals.  From  the  potas- 
sium platinic  chloride  (K2PtCl8)  the  potassium  is  readily  calcu- 
lated. 

ESTIMATION  OF  THE  NITRIC  ACID. 
Fuse  in  a  platinum  crucible,  at  a  low  temperature,  some 
potassium  nitrate,  and  pour  out  on  a  warm  porcelain  slab. 
Ignite  two  or  three  grams  of  silica  in  the  crucible  and  mix 
with  it  about  one-fourth  its  weight  of  the  fused  and  powdered 
nitre.  Bring  the  mixture  to  a  low  red  heat  and  keep  at  that 
temperature  until  it  ceases  to  lose  weight  The  loss  indicates 
the  amount  of  N2O5  from  which  the  HNO3  can  be  calculated. 
Sulphates  and  chlorides  are  not  decomposed  at  this  tempera- 
ture. 


GRAVIMETRIC   ESTIMATION.  101 

CALCIUM  CARBONATE  (CaCO3). 
Estimation  of  the  Caking  -  \ J-  / 

Dissolve  a  weighed  quantity  of  ,trje  carbpnate,  in.  .dilute 
hydrochloric  acid,  avoiding  excess  'of  ,4ei$,,  i(Jd' ^imokiurh'^ 
hydrate  until  slightly  alkaline,  then  ammonium  oxalate,  and 
set  aside  for  twelve  hours.  Decant  the  clear  supernatant 
liquid  to  a  filter,  wash  once  or  twice  by  decantation,  transfer 
the  precipitate  to  the  filter  and  wash  with  small  portions  of 
water,  allowing  each  portion  to  completely  run  through  before 
adding  more.  Dry  and  ignite  the  precipitate,  using  a  strong 
Bunsen  burner  with  a  chimney,  so  as  to  bring  the  crucible  to 
a  white  heat,  until  it  ceases  to  lose  weight ;  if  this  cannot  be 
effected  with  a  Bunsen  burner,  use  a  blast  lamp.  The  residue 
is  weighed  as  calcium  oxide,  CaO. 

ESTIMATION  OF  THE  CARBONIC  ACID. 

This  in  some  cases  is  done  by  ignition,  but  the  method 
adapted  to  all  cases  is  to  treat  the  carbonate  with  an  acid,  as 
HC1  or  H2SO4,  and  estimate  the  loss  of  CO2.  This  is  best 
accomplished  by  using  a  carbonic  acid  apparatus,  that  known 
as  Geissler's  is  probably  the  best. 

An  apparatus  may  be  constructed,  however,  by  taking  a 
small,  wide-mouth  flask,  and  adapting  to  it  a  small  tube  con- 
taining calcium  chloride.  A  weighed  quantity  of  the  material 
is  then  placed  in  the  flask  with  some  water,  a  test  tube  con- 
taining sulphuric  acid  is  put  in  so  as  to  retain  the  acid  until 
the  cork  with  the  drying  tube  is  fixed  in  place,  and  the  whole 
apparatus  weighed. 

The  flask  is  then  turned  over  so  as  to  allow  some  of  the 
acid  to  flow  out,  and  this  is  repeated,  allowing  only  a  small 
quantity  to  run  out  at  a  time,  until  the  carbonate  has  been 
completely  decomposed.  The  flask  is  then  heated,  to  drive 
out  all  the  CO2,  and  after  cooling  the  apparatus  is  weighed. 
The  loss  indicates  the  amount  of  CO2  in  the  material  used. 


102  QUANTITATIVE   ANALYSIS. 

.  .    .  SECTION  II. 

VOLUMETRIC    ESTIMATION. 


Volumetric  analysis  is  the  process  of  determining  the 
amount  of  a  substance  by  the  use  of  a  suitable  reagent  in 
solution  of  a  known  strength.  The  requisite  apparatus  con- 
sists of  a  liter  flask  and  a  1000  cubic  centimeter  cylinder,  in 
which  to  prepare  the  solutions.  A  100  c.c.  burette  (Geissler's 
is  the  most  desirable)  with  a  suitable  holder.  A  few  pipettes, 
assorted  sizes,  some  of  which  are  graduated  to  TV  c.c.  The 
simplest  process  is  that  by  neutralisation,  where  an  alkali  is 
determined  by  neutralizing  it  with  standard  solution  of  oxalic 
acid.  In  such  a  case  it  is  necessary  to  have  a  means  of  know- 
ing when  the  alkali  is  exactly  neutralized.  Such  a  substance 
is  known  as  an  indicator.  Solution  of  litmus  may  be  used,  but 
a  solution  of  phenolphtalein  gives  most  satisfactory  results  ; 
it  is  colorless  in  neutral  or  acid  and  bright  red  in  alkaline 
solution.  It  is  made  by  dissolving  I  part  phenolphtalein  in 
25  parts  alcohol  and  adding  sufficient  water  to  make  100 
parts. 

A  normal  solution  contains  in  every  liter  the  molecular 
weight  of  the  compound  in  grams.  This  is  sometimes 
made  semi-normal,  when  the  compound,  an  acid  for  instance, 
is  dibasic. 

A  deci-normal  solution  is  one-tenth  the  strength  of  the 
normal  one.  While  in  some  of  the  following  examples  the 
normal  solution  is  used,  that  of  deci-normal  strength  is  recom- 
mended for  the  use  of  students. 

VOLUMETRIC  SOLUTION  OF  OXALIC  ACID 
(H2C2O42H2O  —  126). 

The  semi-normal  solution  is  used,  containing  63  grams  in  i  liter. 

Take  of  oxalic  acid,  carefully  purified  by  crystallization, 
63  grams,  transfer  it  to  a  liter  flask,  add  about  800  c.c.  of 


VOLUMETRIC    ESTIMATION.  103 

water,  agitate   until   dissolved  and  bring  the    measure,  with 
water,  to  1000  c.c. 

ESTIMATION  OF  POTASSIUM  'HYDRATE. 

Dissolve  a  convenient  quantity  of  potassium  hydrate,  for 
instance  .500  gram,  in  a  beaker  containing  100  c.c.  of  water  ; 
add  three  or  four  drops  of  the  phenolphtalein  solution  ;  set 
the  beaker  on  a  piece  of  white  paper  and  bring  over  it  a 
burette  containing  some  of  the  oxalic  acid  solution.  Allow 
the  acid  solution  to  run  in  slowly,  constantly  stirring  with  a 
glass  rod,  until  the  red  color  just  ceases  to  be  visible.  The 
proof  of  this  is  that  one  drop  of  a  similar  alkali  solution  will 
restore  the  red  color.  If  the  potassium  hydrate  were  pure,  it 
would  be  found  that  8.94  c.c.  were  used,  but  if  we  suppose 
that  8.50  c.c.  were  necessary,  we  have  every  c.c.  of  the  oxalic 
solution  =  .056  gram  of  KOH,  therefore  8.50  X  .056  =  476 
gram — the  amount  of  KOH  in  the  .500  gram  —  95.20  per 
cent. 

One  c.c.  containing  0.063  gram  of  oxalic  acid  is  the  equiva- 
lent of— 

GRAM. 
Ammonia,  NH3      0.0170 

Ammonium  Carbonate,  NH4HCO3NH4NH2CO2    ...'...  0.0523 

Potassium  Acetate,  KC2H3O2  (after  ignition) 0.0980 

Potassium  Bicarbonate,  KHCO3 o.iooo 

Potassium  Carbonate,  (dry)  K2CO3 0.0690 

Potassium  Citrate,  K3C6H5O7H2O  (after  ignition) 0.1080 

Potassium  Hydrate,  KOH 0.0560 

Potassium  Permanganate,  K2Mn2O8 0.0314 

Potassium  Sodium  Tartrate,  KNaC4H4O64H2O  (after  ignition)  0.1410 
Potassium  Tartrate,  (K2C4H4O6)2H2O  (after  ignition)     .    .    .0.1175 

Lead  Acetate,  Pb(C2H3O2)23H2O 0.1892 

Lead  Subacetate,   Pb2O(C2H3O2)2 0.1367 

Sodium  Bicarbonate,  NaHCO3 0.0840 

Sodium  Borate,  Na2B4O7ioH2O 0.1910 

Sodium  Carbonate,  crystallized,  Na2CO3ioH2O 0.1430 

Sodium  Carbonate,  anhydrous,  Na2CO3 0.0530 

Sodium  Hydrate,  NaHO .  0.0400 

When  carbonates  are  estimated  it  is  necessary  to  boil, 
toward  the  end  of  the  reaction,  after  each  addition '  of  the 
acid,  in  order  to  drive  off  the  CO2,  which  affects  the  indi- 
cator. 


104  QUANTITATIVE   ANALYSIS. 

VOLUMETRIC  SOLUTION  OF   SODIUM    HYDRATE 
(NaOH  =  40). 

The  normal  solution  is  used>  containing  4.0  grams  in  i  liter. 

As  sodium  hydrate  is  never,  or  very  rarely  absolutely  pure, 
it  is  necessary  to  standardize  this  solution.  For  this  take 
something  more  than  the  theoretical  amount  (about  50  grams) 
and  dissolve  in  a  liter  of  water.  Place  100  c.c.  of  the  standard 
oxalic  acid  solution  in  a  beaker,  and,  having  added  the  indi- 
cator, bring  it  under  a  burette  containing  some  of  the  soda 
solution,  and  note  the  number  of  c.c.  necessary  to  exactly 
neutralize  the  acid  solution.  Take  of  the  alkaline  solution 
ten  times  the  number  of  c.c.  necessary  to  neutralize  the  100  c.c. 
of  the  acid  solution,  and  add  sufficient  water  to  bring  the 
measure  to  1000  c.c.  For  instance,  if  it  required  95  c.c.  of 
the  alkaline  solution  to  neutralize  the  100  c.c.  of  the  oxalic 
acid  solution,  then  95  c.c.  X  10  =  950  c.c.  which  amount 
diluted  to  1000  c.c.  would  make  it  exactly  equal  in  strength  to 
the  oxalic  acid  solution. 

One  c.c.  containing  0.040  of  sodium  hydrate  NaOH,  is  the 
equivalent  of — 

GRAM. 

Acetic  Acid,  absolute,  HC2H3O2 0.0600 

Citric  Acid, "crystallized,  H3C6H5O7H2O 0.0700 

Hydrobromic  Acid,  absolute,  HBr 0.0808 

Hydrochloric  Acid,  absolute,  HC1 0.0364 

Hydriodic  Acid,  absolute,  HI 0.1276 

Lactic  Acid,  absolute,  HC3H5O3 0.0900 

Nitric  Acid,  absolute,  HNO3 0.0630 

Oxalic  Acid,  crystallized,  H2C2O42H2O 0.0630 

Sulphuric  Acid,  absolute,  H2SO4 0.0490 

Tartaric  Acid,  crystallized,  H2C4H4O6 0.0750 

VOLUMETRIC   SOLUTION  OF  POTASSIUM    BI- 
CHROMATE (K2Cr2O7=:  294.8). 

A  viginti-normal  solution  is  used,  containing  14.74.  grams  to 

i  liter. 

Place  the  necessary  amount  of  the  salt  (14.74  grams)  in  a 
liter  flask,  add  about  800  c.c  of  water,  agitate  until  dissolved, 
and  bring  the  measure  with  water  to  1000  c.c.  This  solution, 
when  acidified  with  H2SO4,  is  used  in  the  estimation  of  iron  in 


VOLUMETRIC    ESTIMATION.  105 

the  ferrous  condition.  The  end  of  the  reaction  is  determined 
by  taking  out  a  drop  of  the  iron  solution,  and  testing  on  a 
white  porcelain  surface  with  a  drop  of  potassium  ferricyanide 
solution  ;  when  this  ceases  to  give  a  blue  color  the  reaction  is 
at  an  end. 

One  c.c.  containing  0.01474  gram  of  potassium  bichromate, 
K2Cr2O7,  is  the  equivalent  of — 

GRAM. 

Iron  in  ferrous  condition,  Fe 0.01677 

Ferrous  Carbonate,  FeCO3 0.03477 

Ferrous  Sulphate,  FeSO47H2O    .    .    . 0.08337 

Ferrous  Sulphate,  (dry)  FeSO4H2O 0.05097 


VOLUMETRIC  SOLUTION  OF  IODINE  (I  =  126.6). 
A  deci-normal  solution  is  used,  containing  12.66  grams  in  I  liter. 

Weigh  the  necessary  amount  of  iodine  (12.66  grams)  in  a 
stoppered  tube,  to  prevent  loss,  as  well  as  the  corrosive  action 
of  the  fumes  on  the  balance.  Also  weigh  18  grams  of  potas- 
sium iodide  and  place  in  a  liter  flask  with  the  iodine.  Add 
about  200  c.c.  of  water.  The  iodine  dissolves  more  readily  in 
this  strength  of  potassium  iodide  solution,  besides  it  admits 
of  more  thorough  agitation.  When  the  solution  is  complete 
add  water  until  the  liquid  measures  1000  c.c.  Starch  solution 
is  used  as  an  indicator  in  the  determinations  with  the  iodine 
solution. 

One  c.c.  containing  0.01266  iodine  is  the  equivalent  of — 

GRAM. 

Arsenious  Oxide,  As2O3 0.00494 

Potassium  Sulphite,  crystallized,  K2SO32H2O 0.0097 

Sodium  Bisulphite,  NaHSO3 0.0052 

Sodium  Hyposulphite,  crystallized,  Na2S2O35H2O 0.0248 

Sodium  Sulphite,  crystallized,  Na2SO37H2O 0.0126 

Sulphurous  Oxide,  SO2 0.0032 

VOLUMETRIC   SOLUTION    OF  SODIUM  HYPO- 
SULPHITE (Na2S2O35H2O  =  248). 
A  deci-normal  solution  is  used,  containing  24.8  grams  in  I  liter. 

Sodium  hyposulphite  cannot  be  prepared  sufficiently  pure 
to  be  relied  on,  consequently  this  solution  must  be  standard- 
ized, therefore  more  than  the  theoretical  amount  is  taken. 


106  QUANTITATIVE   ANALYSIS. 

The  U.  S.  Pharmacopoeia  recommends  32  grams.  Dissolve 
this  amount  in  1000  c.c.  of  water ;  place  100  c.c.  of  the  standard 
solution  of  iodine  in  a  beaker  and  run  in  the  hyposulphite 
solution  until  the  color  of  iodine  nearly  disappears,  then  add 
a  small  quantity  of  starch  solution  and  continue  until  the  blue 
color  is  discharged.  Multiply  the  number  of  cubic  centi- 
meters of  the  hyposulphite  solution  used  by  10,  and  to  that 
amount  add  sufficient  water  to  bring  the  measure  to  1000  c.c. 
The  substances  estimated  by  this  solution  either  contain  free 
iodine,  or  develop  it  on  the  addition  of  potassium  iodide,  so 
that  starch  solution  may  be  used  as  an  indicator. 

One  c.c.  containing  0.0248  gram  of  sodium  hyposulphite  is 
the  equivalent  of — 

GRAM. 

Bromine,  Br 0.00798 

Chlorine,  Cl 0.00354 

Iodine,  I 0.01266 

VOLUMETRIC    SOLUTION    OF    SILVER    NITRATE 
(AgNO,=  169.7). 

A  deci-normal  solution  is  used,  containing  16.97  grams  in  i  liter. 

As  silver  nitrate  can  be  obtained  or  prepared  perfectly  pure, 
the  necessary  amount  (16.97  grams)  is  dissolved  in  sufficient 
distilled  water  to  make  1000  c.c. 

In  testing  one  of  the  following  compounds,  ammonium 
chloride  for  instance,  a  weighed  amount  is  taken,  dissolved 
in  water,  and  a  few  drops  of  potassium  bichromate  solution 
added.  The  silver  nitrate  solution  is  then  run  in  until  a  red 
precipitate  remains  permanently.  The  silver  combines  with 
the  chlorine  until  the  latter  is  all  used,  when  it  forms  with 
the  chromic  acid  red  silver  chromate,  so  that  its  formation 
indicates  the  end  of  the  reaction. 

One  c.c.  containing  0.01697  gram  of  nitrate  silver  is  the 
equivalent  of — 

GRAM. 

Ammonium  Bromide,  NH4Br 0.00978 

Ammonium  Chloride,  NH4C1 0.00534 

Ferrous  Bromide,  FeBr2 0.01077 

Ferrous  Iodide,  FeI2 0.01545 

Hydrocyanic  Acid,  absolute,  HCN 0.00270 


VOLUMETRIC    ESTIMATION.  107 

GRAM. 

Hydriodic  Acid,  HI 0.01276 

Potassium  Bromide,  KBr 0.01188 

Potassium  Chloride,  KC1 .  0.00744 

Potassium  Cyanide,  KCN • 0.01300 

Sodium  Bromide,  NaBr 0.01028 

Sodium  Chloride,  NCI 0.00584 


TABLE   OF   ELEMENTS. 

ELEMENTS.                                                                          SYMBOLS.  ATOMIC    WEIGHTS. 

Aluminium Al 27 

Antimony Sb 120 

Arsenic As 74.9 

Barium Ba 136.8 

Beryllium Be 9 

Bismuth Bi 210 

Boron B n 

Bromine      Br 79.8 

Cadmium Cd in.8 

Caesium Cs     . 132-6 

Calcium Ca 40 

Carbon C 12 

Cerium Ce 140.2 

Chlorine      Cl 35.4 

Chromium Cr 52.4 

Cobalt Co 58.9 

Copper Cu 63.2 

Didymium Di !42-3 

Erbium .  E I^>S-9 

Fluorine      Fl 19 

Gallium G      68.8 

Gold Au 196.2 

Hydrogen H I 

Indium :  In U3-4 

Iodine I 126.6 

Iridium Ir      l92-7 

Iron     .    .    .    .  ' Fe 55.9 

Lanthanum La 138.2 

Lead Pb 206.5 

Lithium Li 7 

Magnesium Mg .    .  24 

Manganese Mn 54 

Mercury Hg l99-7 

Molybdenum      Mo 95.5 

Nickel Ni 58 

Niobium Nb 94 

Nitrogen N ,    .    .    .  14 


108  QUANTITATIVE   ANALYSIS. 

TABLE  OF  ELEMENTS.— Continued. 

ELEMENTS.                                                                            SYMBOLS.  ATOMIC    WEIGHTS. 

Osmium Os 198.5 

Oxygen O      1 6 

Palladium Pd 105.7 

Phosphorus P 31 

Platinum Pt 194-4 

Potassium K 39 

Rhodium Rh 104.1 

Rubidium Rb 85.3 

Ruthenium Ru 104.2 

Scandium Sc 44 

Selenium Se 78.8 

Silicon Si 28 

Silver Ag 107.7 

Sodium Na 23 

Strontium Sr 87.4 

Sulphur S 32 

Tantalum Ta 182 

Tellurium Te 128 

Thallium Tl 203.7 

Thorium Th 233 

Tin Sn II7-7 

Titanium Ti 48 

Tungsten W 183.6 

Uranium U      23&>5 

Vanadium V $1.3 

Ytterbium Yb    . 172.7 

Yttrium Y      89.8 

Zinc Zn 64.9 

Zirconium Zr 90 

Davyum,  Decipium,  Germanium,  Neptunium  and  a  few  others,  are  hardly 
sufficiently  known  to  warrant  their  being  placed  in  the  above  list. 


INDEX. 


PAGE 

Acetates,  detection  of.  .........................  81,  83 

reactions  of.  ..................................  77 

Acetic  acid  ............................................  77 

Acid,  acetic  ..........................................  77 

benzoic  .........................................  80 

boric  ............................................  74 

carbolic  .......................................  79 

carbonic  .......................................  74 

estimation  of.  .............................  101 

chloric  .........................................  70 

citric  ............................................  79 

gallic  ............................................  80 

hydriodic  ......................................  68 

hydrobromic  .................................  67 

hydrochloric  .............................  20,  67 

hydrocyanic  .................................  68 

hydroferricyanic  ...........................  75 

hydroferrocyanic  ...........................  75 

hydrofluoric  .................................  68 

hydrosulphuric  ..............................  70 

hypochlorous  ...............................  70 

hypophosphorous  ..........................  72 

hyppsulphurous  .............................  71 

lactic  ............................................  77 

malic  ................  ...........................  78 

metaphosphoric  .............................  73 

nitric  .........................................  72 

estimation  of  ..............................  100 

oleic  ..............  .  ..............................  77 

orthophosphoric  ...........................  72 

oxalic  ..........................................  78 

phosphoric  ...................................  72 

pyrophosphoric  .....................  .......  73 

salicylic  .......................................  80 

silicic    .........................................  74 

stearic  ..........................................  77 

succinic  .......................................  78 

sulphuric  .....................................  71 

estimation  of.  .............................  100 

sulphurous  ....................................  71 

tannic  ..........................................  80 

tartaric  ........................................  78 

thiosulphuric  .................................  71 

valerianic  .........................  ,  ...........  77 

Acids,  analytical  reactions  of.  .................  67 

detection  of.  ...............................  82-87 

Acidum  hydrochloricum  ..........................  21 

Alcohol,  reactions  and  tests  of.  .................  89 

amylic,  detection  of.  ......................  89 

Alkaloids,  reactions  and  tests  of  ...........  91 

Aluminium,  analytical  reactions  of.  .........  47 

hydrate  ........................................  29 

Ammonia,  preparation  of.  .......................  23 

properties  of  .................................  23 

Ammonium,  analytical  reactions  of.  ........  34 

nitrate  .........................................  27 

Analysis  of  Acids  — 

Group  I  ..........................  .  .............  69 

"     II  .......................................  75 

"  HI  .......................................  81 

Bases  — 

Group  1  .......................................  35 

"     II  .......................................  38 


VII 


PACK 

Analysis  of  soluble  salts 83 

insoluble  salts 86 

Antimony,  analytical  reactions  of 53 

Marsh's  test  for 53 

Reinsch's  test  for 53 

Aqua  chlori 20 

Arsenates,  analytical  reactions  of 51 

detection  of. 85 

Arsenic,  analytical  reactions  of 50 

Fleitman's  test  for 52 

Marsh's  test  for 51 

Reinsch's  test  for 51 

Arsenious  oxide 50 

Atomic  weights,  list  of. 107 

Atropine 93 

Balance,  rules  for  using 97 

Barium,  analytical  reactions  of. 35 

chloride,  estimation  of. 98 

Benzoic  acid 80 

Bismuth,  analytical  reactions  of. 58 

Bleaching 20 

Boric  acid,  analytical  reactions  of 74 

Bromides,  detection  of. 69,  84 

Cadmium,  analytical  reactions  of. 59 

Caffeine 93 

Calcium,  analytical  reactions  of. 37 

carbonate,  estimation  of. 101 

phosphate,  preparation  of. 27 

Cane  sugar 91 

Carbolic  acid 79 

Carbon  dioxide,  preparation  of. 24 

Carbonic  acid,  analytical  reactions  of. 74 

estimation  of. 101 

Cerium,  analytical  reactions  of 46 

Chloral  hydrate,  reactions  and  tests  of.....  89 

Chloric  acid 70 

Chlorides,  detection  of. 69,84 

Chlorine,  estimation  of. 99 

preparation  of 19 

water 20 

Chloroform,  reactions  and  tests  of. 88 

Chromium,  analytical  reactions  of. 47 

Cinchonidine,  reactions  and  tests  of. 92 

Cinchonine,  reactions  and  tests  of. 93 

Citric  acid,  analytical  reactions  of. 79 

detection  of. 81,  84 

Cobalt,  analytical  reactions  of. 41 

Copper,  analytical  reactions  of 58 

estimation  of. 99 

sulphate,  estimation  of. 99 

preparation  of. 30 

Cyanides,  detection  of. 69,  85 

Deci-normal  solution 102 

Dextrose 90 

Drying 98 

Elements,  table  of. 107 

Gallic  acid,  analytical  reactions  of. 80 

detection  of. 83 

Glauber  salt 21 

Glucose,  reactions  and  tests  of 90 

Glycerin,  reactions  and  tests  of. 90 

Gold,  analytical  reactions  of. 54 

Grape  sugar go 

Gravimetric  estimation 97 

Group  I,  Acids 67 

"     II,     "     70 

"  HI,     "     77 

"      I,  Bases 33 


109 


110 


INDEX. 


PAGE 

Group  II,  Bases 35 

III,  "     40 

IV,  "       : 45 

V,      "     50 

VI,      "     57 

VII,      "     62 

Group  reagents 66 

Hydrates,  analytical  reactions  of 70 

Hydriodic  acid,  analytical  reactions  of 68 

Hydrobromic  acid,  analytical  reactions  of  67 

Hydrochloric  acid,  analytical  reactions  of..  67 

preparation  of 20 

properties  of. 21 

Hydrocyanic  ecid,  analytical  reactions  of.  68 

Hydroferricvanic  acid 75 

Hydroferrocyanic  acid 75 

Hydrofluoric  acid 68 

Hydrogen 17 

preparation  of 17 

properties  of. 17 

arsenide 51 

Hydrosulphuric  acid 70 

Hypochlorous  acid 70 

Hypophosphorous  acid 72 

Hype-sulphurous  acid 71 

Ignition 98 

Indicator 102 

Iodides,  detection  of 69,85 

lodoform,  reactions  and  tests  of 88 

Iron,  analytical  reactions  of. 45 

Lactic  acid 77 

Lead,  acetate  of. 30 

analytical  reactions  of 63 

Lithium,  analytical  reactions  of 34 

Magnesium,  analytical  reactions  of. 37 

carbonate 28 

oxide 29 

sulphate 28 

Malic  acid 78 

Manganese,  analytical  reactions  of. 40 

Marsh's  test  for  arsenic 51 

Mercuric  compounds 57 

analytical  reactions  of. 57 

Mercurous  compounds 63 

analytical  reactions  of 63 

Metaphosphoric  acid 73 

Morphine,  reactions  and  tests  of 91 

Neutralization 102 

Nickel,  analytical  reactions  of. 42 

Nitrates,  detection  of 85 

Nitric  acid,  analytical  reactions  of  72 

preparation  of 24 

properties  of. 24 

Nitrogen,  preparation  of. 22 

properties  of 23 

Normal  solution 102 

Oleic  acid 77 

Oxidation -. 20 

Oxides 70 

Oxygen,  preparation  of 21 

properties  of 22 

Phenolphtalein 102 

Phosphates,  detection  of. 85 

Phosphoric  acid,  analytical  reactions  of...  72 

Platinum,  analytical  reactions  of. 54 

Potassium,  analytical  reactions  of. 33 

chloride,  preparation  of. 26 

estimation  of 100 


PAGE 

Potassium  hydrate,  estimation  of. 103 

and  sodium  tartrate 26 

Precautions  on  Group  II,  Bases 

"      IV,      »     

V,      "     

"     VI,      "     

"    VII,      "     

Precipitation 

Purple  of  Cassius 54 

Quantitative  analysis,  preliminary  direc- 
tions   97 

Quinine,  reactions  and  tests  of. 92 

Reinsch's  test  for  arsenic 51 

Rochelle  salt 26 

Saccharose,  reactions  and  tests  of. 91 

Salicin 94 

Salicylic  acid 80 

Santonin,  reactions  and  tests  of. 94 

Seminormal  solution 102 

Silicic  acid ...  74 

Silver,  analytical  reactions  of. 62 

Sodium,  analytical  reactions  of 33 

sulphate 21 

Solubilities,  chart  of. 87 

Stannic  compounds 54 

Stannous  compounds 53 

Starch 91 

Stearic  acid 77 

Strontium,  analytical  reactions  of 36 

Strychnine,  reactions  and  tests  of 92 

Succinic  acid 78 

Sulphates,  detection  of. 76,  84 

Sulphuretted  hydrogen 70 

Sulphuric  acid,  analytical  reactions  of. 71 

Sulphurous  acid,  analytical  reactions  of...  71 

Summary  Group  I,  Acids 69 


II, 

III,      "     . 

I,  Bases. 

II, 

III, 

IV, 

V, 

VI, 

VII, 


55 

6o 

64 

Symbols,  table  of. 107 

Tannates,  detection  of 85 

Tannic  acid,  analytical  reactions  of. 80 

Tartaric  acid,  analytical  reactions  of 78 

Tartrates,  detection  of 81,84 

Thalleioquin 92 

Thiosulphuric  acid 71 

Tin,  analytical  reactions  of. 53 

Valerianic  acid 77 

Veratrine,  reactions  and  tests  of. 93 

Volumetric  analysis 102 

estimation  of  alkalies 103 

solutions 102 

iodine 105 

oxalic  acid 102 

potassium  bichromate 104 

silver  nitrate 106 

sodium  hydrate 104 

sodium  hyposulphite 105 

Wash  bottle 28 

Water 70 

Zinc,  analytical  reactions  of 41 

sulphate,  preparation  of..... 19 


14  DAY  USE 

RETURN  TO  DESK  FROM  WHICH  BORROWED 

LOAN  DEPT. 


f  This  book  is  due  on  the  last  date  stamped  below,  or 
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Renewed  books  are  subject  to  immediate  recall. 

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